Category: CBSE Notes

By going through these CBSE Class 12 Chemistry Notes Chapter 13 Amines, students can recall all the concepts quickly.

Amines Notes Class 12 Chemistry Chapter 13

Amines: Amines can be considered as derivatives of ammonia, obtained by replacement of H of NH₃ by alkyl /aryl group.

→ Structure of Amines: Like ammonia, the N atom of amines is trivalent and Carnes an unshared pair of electrons. N orbitals in amines are sp3 hybridized and the geometry of amines is pyramidal.

(Pyramidal shape of trimethylamine)

→ Classification of Amines: Amines are classified as primary (1° secondary (2°) and tertiary (3°) depending upon the no. of hydrogen atoms replaced by alkyl/aryl groups in ammonia molecule.

Amines are simple if R = R’ = R”
They are termed as mixed when R, R’, R” are different.

→ Nomenclature of Amines: In the IUPAC system ‘e’ of the alkane is replaced by ‘amine’ like Alkanamine. The simplest arylamine is aniline (C₆H₅—NH₂) or benzene amine

→ Preparation of Amines:
1. Reduction of Nitro Compounds

2. Ammonolysis of Alkyl Halides

The order of reactivity of halides is RI > RBr > RCI.

3. Reduction of Nitriles

4. Reduction of amides

5. Gabriel Phthalimide Synthesis: This method is used to prepare pure primary amines. Aromatic amines cannot be prepared by this method.

Hoffman Bromamide Degradation Reaction: Primary amines are obtained when an amide is treated with Br₂ in an aqueous solution of NaOH.

Physical Properties:

1. Lower aliphatic amines are gases with a fishy odour. Primary amines with three or more C atoms are liquids and still higher ones are solids.
2. Lower aliphatic amines are soluble in water because they can form H-bonds with water molecules. Higher amines are essentially insoluble in water.
3. The order of boiling points of isomeric amines is Primary > Secondary > Tertiary. It is due to the presence of intermolecular H-bonding which is more in primary amines (due to the presence of two H- atoms attached to N than only one in secondary amines), less in 2° amines and absence in tertiary amines.
   

Chemical Reactions:
1. Basic character of Amines: Amines being basic in character (Lewis bases) react with acids to form salts.

The basic character of amines can be better understood in terms of their P^K_b and k_b values as explained below:


The larger the value of K_b or lower the value of P^K_b, the stronger is the base. P^K_b value of ammonia is 4.75. Aliphatic amines are stronger bases than ammonia due to the + I-effect of alkyl groups. The availability of lone pair of electrons on N increases. On the other hand, aromatic amines are weaker bases than ammonia due to the electron-withdrawing nature of the aryl group. The P^K_b values of few amines are given below:

Table: P^K_b values of Amines in Aqueous Phase:

Name of amine	PKb
Methenamine	3.38
N-Methylmethanamine	3.27
N, N-Dimethylmethanamine	4.22
Ethanamine	3.29
N-Ehtylethamine	3.25
Benzenamine	9.38
Phenylmethanamine	4.70
N-Methylaniline	9.30
N, N-Dimethylaniline	8.92

Besides the + I or – I effect, the interplay of other factors like solvation effect, steric hindrance etc. affect the basic strength of amines.

→ Structure-Basicity Relationship of Amines: The basicity of amines is related to their structure. The basic character of an amine depends upon the ease of formation of the cation by accepting a proton from the acid.

The more stable the cation is relative to the amine, the more basic is the amine.
(a) Alkamines versus Ammonia
Let us consider the reaction of an alkanamine and ammonia with a proton to compare their basicity.

Due to the electron releasing nature of the alkyl group, it (R) pushes electrons towards nitrogen and thus makes the unshared electron pair more available for sharing with the proton. Moreover, the substituted ammonium ion formed from the amine gets stabilised due to dispersal of the positive charge by the +1 effect of the alkyl group. Hence, alkylamines are stronger bases than ammonia.

Thus the basic nature of aliphatic amines should increase with the increase in the number of alkyl groups. This trend is followed in the gaseous, phase. The order of basicity of amines in the gaseous phase follows the expected order: tertiary amine > secondary amine > primary amine > NH₃.

The trend is not regular in the aqueous state as evident by their P^K_b values given in Table. In the aqueous phase, the substituted ammonium cations get stabilised not only by the electron releasing effect of the alkyl group (+1) but also by hydrogen bonds and solvation with water molecules. The greater the size of the ion, the lesser will be the solvation and the less stabilised the ion. The order of stability of ions are as follows:

Decreasing order of extent of H-bonding in water and stabilization by solvation: Greater the stability of the substituted ammonium cation, the stronger base is the corresponding amine. Thus the order of basicity of aliphatic amines should be primary > secondary > tertiary, which is opposite to the inductive effect based order. Secondly, when the alkyl group is small like – CH₃ group, there is no steric hindrance to H-bonding.

In case the alkyl group is bigger than the CH₃ group, there will be some steric hindrance to H-bonding. Therefore, the change of nature of the alkyl group, e.g., from CH₃ to – C₂H₅ results in a change in the order of basic strength. Thus, there is a subtle interplay of the inductive effect, solvation effect and steric hindrance of the alkyl group which decides the basic strength of alkylamines in the aqueous state.

The order of basic strength in the case of methyl-substituted amines and ethyl substituted amines in an aqueous solution is as follows:
(C₂H₅)₂ NH > (C₂H₅)₃N > C₂H₅NH₂ > NH₃
(CH₃)₂NH > CH₃NH₂ > (CH₃)₃N > NH₃

(b) Arylamine Vs Ammonia: High value of P^K_b of aniline suggests that it is a much weaker base than aliphatic amines or even ammonia. It is due to the resonance shown by aniline.

Aniline is a resonance hybrid of 5 canonical structures out of which there is no lone pair of electrons available for sharing on N in II, III and IV structures. It results in the unshared electron pair on N being in conjugation with the benzene ring making it less available for protonation.

On the other hand, aclidinium ion obtained by accepting a proton can have only two resonating structures (Kekule) as shown below:

The greater the no, of resonating structures, the greater is the stability. Thus aniline is more resonance stabilized than anilinium ion. Hence the proton acceptability or the basic nature of aniline and other aromatic amines would be less than ammonia.

In the case of substituted aniline, electron-releasing groups like – OCH₃ – CH₃ increase the basic strength whereas electron-withdrawing groups like – NO₂, – SO₃H, – COOH, – X decrease it.
Thus (C₂H₅)₂ NH > C₂H₅NH₂ > NH₃ > C₆H₅NH₂

2. Alkylation: Amines undergo alkylation with alkyl halides.

3. Acylation: Aliphatic and aromatic primary and secondary amines react with acid chlorides, anhydrides and esters by a nucleophilic substitution reaction. This reaction is known as acylation.

It is the replacement of H of – NH₂ group
or

group by the acyl group.


The reaction of amines with benzoyl chloride (CH₅ COCl) is called Benzoylation.

4. Carbylamine reaction: It is a test of primary amines-both aliphatic and aromatic.

5. Reaction with Nitrous acid: All the three classes of amines react differently with nitrous acid, HNO₂ (prepared in situ from HCl + NaNO₂)


Secondary and tertiary amines react differently.

6. Reaction with aryl sulphonyl chloride: Benzensulphonyl chloride (C₆H₅SO₂Cl) called Hinsberg’s reagent reacts with 1° and.2° amines.

(c) Tertiary amines (R₃N) do not react with benzene sulphonyl chloride. This property of 1°, 2° and 3° amines is made use of in their distinction and separation Now in place of benzene sulphonyl chloride,

p-toluenesulphonyl chloride is used.

7. Electrophilic Substitution: NH₂ group fused in the ring is a powerful activating group and is ortho and para directing.
(a) Bromination:

Protection of the highly activating -NH₂ group can be done by acetylation with acetic anhydride in which case only monobromo substituted aniline is the product.

The lone pair of electrons on N of acetanilide interacts with oxygen atom due to resonance as shown below:

Hence the lone pair of electrons on nitrogen is less readily available for donation to the benzene ring by resonance. Thus the activating effect of – NHCOCH₃ is less than that of the – NH₂ group.

(b) Nitration: Direct nitration of aniline yields tarry oxidation products in addition to nitro derivatives.

However, by protecting the amino group by acetylation reaction with acetic anhydride the nitration can be controlled and the para nitro derivative obtained as the major product.

(c) Sulphonation:

Aniline does not undergo Friedel Crafts reaction (alkylation and acetylation) due to the salt formation with AlCl₃– the Lewis acid-which is used as a catalyst. The N of -NH₂ acquires a positive charge and hence acts as a strong deactivating group for further reaction.

Diazonium Salts are of the type Ar N₂⁺ X^– where Ar stands for an aryl group and X^– can be Cl^–, Br, HSO₄^–, BF₄^– etc. N₂⁺ group is called the diazonium group.

C₆H₅ N₂⁺Cl^– is called benzene diazonium chloride and C₆H₅ N₂⁺ HSO₄^– is known as benzene diazonium hydrogen sulphate. The stability of the arene diazonium ion is explained on the basis of resonance.

Method of Preparation:

Physical Properties:
1. It is a colourless crystalline solid.
2. It is readily soluble in cold water and stable in it but reacts with water when warmed.
3. It decomposes easily in the dry state. Therefore it is used immediately after its preparation and is not stored.

Chemical Reactions A: Reactions involving displacement of Nitrogen -N₂⁺ group is a very good leaving group. Therefore, it is substituted by other groups such as Cl^–, Br^–, I^–, CN^– and OH^– which displace nitrogen from the aromatic ring.
1. Replacement by halide or cyanide ion: Sandmeyer Reaction

Alternatively, chlorine or bromine can also be introduced in the benzene ring by treating diazonium salt solution with corresponding halogen acid in the presence of copper powder. This is referred to as the Gatterman reaction.

The yield in the Sandmeyer reaction is found to be better than the Gattermann reaction.

2. Replacement by Iodide Ion: Iodine is not easily introduced into the benzene ring directly, but, when the diazonium salt solution is treated with potassium iodide, iodobenzene is formed.
Ar N₂⁺ Cl^– + KI → Arl + KCl + N₂

3. Replacement by Fluoride Ion: When arene diazonium chloride is treated with fluoroboric acid, arene diazonium fluoroborate is precipitated which on heating decomposes to yield aryl fluoride.

4. Replacement by H: Certain mild reducing agents like hypophosphorous acid (phosphinic acid) or ethanol reduce diazonium salts to arenes and themselves get oxidised to phosphorous acid and ethanol, respectively.
Ar N₂⁺ Cl^– + H₃PO₂ + H₂O → ArH + N₂ + H₃PO₃ + HCl
Ar N₂⁺ Cl^– + CH₃CH₂OH → ArH + N₂ + CH₃CHO + HCl

5. Replacement by hydroxyl group: If the temperature of the diazonium salt solution is allowed to rise up to 283 K, the salt gets hydrolysed to phenol.
Ar N₂⁺ Cl^– + H₂O → ArOH + N₂ + HCl

6. Replacement by – NO₂ group: When diazonium fluoroborate is heated with aqueous sodium nitrite solution in the presence of copper, the diazonium group is replaced by the NO₂ group.

(B) Reactions involving retention of diazo group: Coupling Reactions
(a) Reaction with Phenol in the presence of a weakly alkaline medium results in the formation of p-hydroxy azobenzene which is an orange dye.

(b) Reaction with aniline: It reacts with aniline in a weakly acidic medium to form p-amino azobenzene (yellow dye)

→ Importance of Diazonium salts in the synthesis of Aromatic Compounds: These salts are very good intermediates used to introduce – F, – Cl, – Br,-I, – CN, – OH, – NOz groups into the benzene ring. Ar-F or Ar-I cannot be prepared by direct halogenation. The cyanobenzene which cannot be prepared by the S_N reaction of chlorobenzene can be easily obtained from diazonium salts. These compounds are useful for preparing several azo dyes by coupling reactions.

By going through these CBSE Class 12 Chemistry Notes Chapter 12 Aldehydes, Ketones and Carboxylic Acids, students can recall all the concepts quickly.

Aldehydes, Ketones and Carboxylic Acids Notes Class 12 Chemistry Chapter 12

→ Aldehydes and Ketones both contain carbonyl C = O group and hydrogen while in Ketones, it is bonded to two carbon atoms.

→ Carbon in both aldehydes and Ketones is sp² hybridized:
If the Carbonyl group Carboxylic acid. is attached to -OH group, it is called Carboxylic Acid


Common names:

IUPAC Names: The IUPAC names of open chain aliphatic aldehydes and ketones are derived from the names of the corresponding alkanes by replacing the ending – e with – al and – one respectively.


Table: Common and IUPAC Names of Some Aldehydes and Ketones:


→ Structures of the Carbonyl group): The carbonyl carbon atom is sp² hybridized and forms three sigmas (σ) bonds.

C = O is polarised due to the higher electronegativity of oxygen relative to carbon. Hence the carbonyl carbon is an electrophilic (Lewis acid) and carbonyl oxygen, a nucleophile (Lewis base) centre.

The carbonyl group is highly polar.

Preparation of Aldehydes and Ketones:
1. Aldehydes are prepared by the partial oxidation of primary alcohols while ketones are obtained from secondary alcoholic.

2. Catalytical dehydrogenation of primary alcohols with red hot Cu gauze at 573 K gives aldehydes and secondary alcohols give ketones.

3. From ozonolysis of alkenes:


4. From Alkynes from hydration:

Preparation of Aldehydes only
1. Rosenmund’s Reaction: Reduction of acid Chlorides with H2 and Pd/BaS04 catalyst in boiling xylene solution.

2. From Nitriles and esters: Stephen Reaction

Esters are reduced to aldehydes with DIBAL-H [diisobutyl- aluminium hydride]

3. From Aromatic hydrocarbons:
(a) With chromyl chloride (Cr02Cl2): Etard’s Reaction

(b) With Chromic Oxide (Cr03)

(c) By side-chain chlorination followed by hydrolysis

(d) By Gatterman-Koch Reaction

Methods of Preparation for Ketones:
1. From acid chlorides

2. From nitriles

3. From benzene: By Friedel-Crafts acylation:

→ Physical Properties:
1. Methanal is a gas at room temperature. Ethanol is a volatile liquid. Other aldehydes and ketones are liquid or solid at room temperature.

2. B. Pts of aldehydes and ketones are higher than hydrocarbons and ethers of comparable molecular masses due to weak molecular association in aldehydes and ketones due to dipole-dipole interactions. Their B.Pts. are lower than those of alcohols due to the absence of H-bonding.

The following compounds of molecular masses 58 and 60 are ranked in order of increasing boiling points.

	B.Pt (K)	M.Mass
n-Butane	273	58
Methoxyethane	281	60
Propanal	322	58
Aoelone	329	58
Propan-1-ol	370	60

3. The lower members of aldehydes and ketones such as methanal, ethanal and propanone are miscible with water in all proportions because they form H-bonds with water.

Solubility of aldehydes and ketones decreases rapidly on increasing the length of the alkyl group.

4. The lower aldehydes have a sharp pungent odour. As molecular mass increases, the odour becomes less pungent and more fragrant.

Chemical Reactions of Aldehydes and Ketones:
1. Nucleophilic addition Reactions: On the attack of a nucleophile on the carbonyl carbon, the hybridization of C changes from sp² to sp³.

Aldehydes are more susceptible to nucleophilic addition reactions than ketones due to steric and electronic factors.

The reactivity of aldehydes and ketones towards nucleophilic addition reactions is

The polarity of the carbonyl group is reduced in benzaldehyde due to resonance and hence it is less reactive than propanal.

1. Addition of Hydrogen cyanide (HCN):
HCN + OH- ⇌ -CN + H2O

2. Addition of Sodium bisulphite (NaHSO3):

3. Addition of Grignard’s reagent: They give 1°, 2°, 3° alcohols.

4. Addition of alcohols: Acetals/Ketals are formed

Ketones, under similar conditions, react with ethylene glycol to form cyclic products known as ethylene glycol ketals.

5. Addition of ammonia and its derivatives [Addition- Elimination].

Z = alkyl, aryl, OH, NH₂, C₆H₅NH, NHCONH₂ etc.


[Table: Some N-Substituted Derivatives of aldehydes and ketones

2, 4-DNP-derivatives are yellow, orange or red solids, useful for characterization of aldehydes and ketones.

2. Reduction:
1. Reduction to alcohols: Aldehydes and ketones are reduced to primary and secondary alcohols respectively by sodium borohydride (NaBH₄) or lithium aluminium hydride (LiAlH₄) as well as by catalytic hydrogenation.

2. Reduction to hydrocarbons: The carbonyl group of aldehydes and ketones is reduced to CH2 group on treatment with zinc-amalgam and concentrated hydrochloric acid [Clemmensen reduction) or with hydrazine followed by heating with sodium or potassium hydroxide in a high boiling solvent such as ethylene glycol (Wolff-Kishner reduction).

3. Oxidation: Aldehydes differ from ketones in their oxidation reactions,

Due to the easy oxidation of aldehydes as compared to ketones, they can be distinguished from the following two steps:
1. Tollen’s test:
R—CHO + 2 [Ag(NH₃)₂]⁺ + 3 OH^– → RCOO^– + Ag (s) + 2 H₂O + 4 NH₃
On warming an aldehyde with freshly prepared ammonical AgNO₃ solution (Tollen’s reagent), a bright silver mirror is produced.

2. Fehling’s Test: On heating an aldehyde with Fehling reagent, a reddish-brown precipitate is obtained.
RCHO + 2 Cu²⁺ + 5 OH^– → RCOO^– + Cu₂O + 3 H₂O. (red-brown ppt.)

→ Oxidation of methyl ketones by haloform reaction: All those carboxyl compounds containing the – COCH₃ group respond to this test.

If NaOI is used, yellow ppt. of CHI₃ is formed which is used to detect the CH₃C = O group.

4. Reactions due to α-hydrogen atom: α-hydrogens in aldehydes and ketones are acidic in nature due to the strong electron-withdrawing effect of the carbonyl group and resonance stabilisation of the conjugate base.

1. Aldol Condensation: Aldehydes and ketones having at least one α-hydrogen undergo Aldol Condensation in the presence of dil. alkali to form β-hydroxy aldehydes (aldol) or β-hydroxy ketones [ketol] respectively.


2. Cross Aldol Condensation: When aldol condensation is carried out between two different aldehydes and/or ketones, it is called Cross aldol condensation.

Ketones can also be used as one component in the cross aldol condensation.

5. Other Reactions:
1. Cannizzaro Reactions: Aldehydes that do not have an a- hydrogen atom, undergo self oxidation-reduction called Disproportionation reactions on treatment with concentrated alkali.

2. Electrophilic substitution reactions: Aromatic aldehydes and ketones undergo electrophilic substitution reactions at the ring and the carbonyl group acts as a deactivating and meta-directing group.

Uses of aldehydes and ketones

• Formalin (aqueous solution of formaldehyde) is used to preserve biological specimen and to prepare bakelite.
• Acetaldehyde is used to prepare acetic acid, ethyl acetate, vinyl acetate, polymers and drugs.
• Benzaldehyde is used in perfumery and in dye industries.
• Acetone and ethyl methyl ketone are common industrial solvents.

II. Carboxylic Acids: Carbon compounds containing a carboxylic -COOH group are called Carboxylic acids.

→ Nomenclature and Structure of Carboxylic Group: In the IUPAC system last – e of alkanes is replaced with – oic acid. For naming compounds containing more than one carboxylic group, the ending – C of the alkane is retained.

Table: Names and Structures of Some Carboxylic Acids:

→ Structure of Carboxylic Group: Because of the possible resonance structure shown below, the carboxylic carbon is less electrophilic. Then carbonyl carbon and the bonds of the carboxylic carbon lie in the plane and are separated by 120°.

Methods of Preparation of Carboxylic Acids:
1. From Primary Alcohols and Aldehydes: Primary

2. From Alkyl benzenes:

3. From Nitriles and amides: By Hydrolysis

2. Due to extensive hydrogen bonding, carboxylic acid molecules get associated and thus have higher boiling points as compared to aldehydes, ketones and even alcohols of comparable molecular masses.

(In vapour state are in an aprotic solvent.)

3. Simple aliphatic carboxylic acids having up to four C atoms are miscible in water the formation of hydrogen bonds with water.

The solubility decrease with an increasing number of carbon atoms.

(Hydrogen bonding of RCOOH with H₂O)

Chemical Reaction:
Reactions involving cleavage of O-HBond

Acidic character Reactions with metals and alkalies:
2RCOOH + 2 Na → 2 RCOO^– Na⁺ + H₂ Sodium carboxylate
RCOOH + NaOH → RCOO^– Na⁺ + H₂O
RCOOH + NaHCO₃ → RCOO^– Na⁺ + H₂O + CO₂

Carboxylic acid dissociates in water to give resonance-stabilized carboxylate anion and the hydronium ion.

For the above reaction:

Where K_eq, is the equilibrium constant and K_a is the acid dissociation constant.
p^ka = – log k_a

The p^ka of hydrochloric acid is – 7.0, whereas p^ka of trifluoracetic acid (tine strongest organic acid), benzoic acid, and acetic acid are 0.23, 4.19 and 4.76 respectively.

Smaller the p^ka, the stronger the acid.
Carboxylic acids are weaker than mineral acids, but they are stronger acids than alcohols and many simple phenols [pka for phenols is ~ 10 and for ethanol ~ 16). Carboxylate anion is more resonance stabilized than phenoxide ion as the – ve charge is spread on two oxygen atoms rather than one in phenoxide ion.

Effect of Substituents on the acidity of carboxylic acids: Electron withdrawing groups increase the acidity of carboxylic acids by stabilizing the conjugate base through delocalisation of the negative charge by inductive and/or resonance effects. Conversely, electron-donating groups decrease the acidity by destabilizing the conjugate base.

The effect of the following groups in increasing acidity order is Ph < I < Br < Cl < F < CN < NO2 < CF3.

Thus the following acids are arranged in order of increasing acidity/based on p^ka values.

Direct attachment of groups such as phenyl or vinyl to the carboxylic acid, increases the acidity of corresponding carboxylic acid, contrary to the decrease expected due to the resonance effect shown below:

This is because of the greater electronegativity of sp² hybridised carbon to which carboxyl carbon is attached. The presence of an electron-withdrawing group on the phenyl of the aromatic carboxylic acid increases their acidity while electron-donating groups decrease their acidity.

Reactions Involving Cleavage of C—OH Bond:
1. Formation of Anhydride

2. Esterification:

Mechanism of Esterification:

3. Reaction involving PCl₅ PCl₃ SOCl₂: Thionyl chloride (SOCl₂) is preferred as the other two products are gases.
RCOOH + PCl₅ → R COCl + POCl₃ + HCl
3RCOOH + PCl₃ → 3 ROCl + H3PO₃
RCOOH + SOCl₂ → RCOCl + SO₂ ↑ + HCl ↑.

4. Reactions with ammonia:

Reactions Involving: COOH Group:
1. Reduction:

2. Decarboxylation:

Kolbe’s Electrolysis

Substitution reactions in the hydrocarbon part:
1. Halogenation: Hell-Volhard-Zelinsky Reaction

2. Ring Substitution: —COOH group present in the ring is a deactivating group and meta-directing group. They do not, however, undergo Friedel-Craft reaction.

It is because the carboxylic group (— COOH) is a deactivating group and the catalyst aluminium chloride (Lewis acid) gets bonded to the carboxylic group.

Uses of Carboxylic acids: Methanoic acid is used in rubber, textile, dyeing, leather and electroplating industries. Ethanoic acid is used as a solvent and as vinegar in the food industry. Hexanedioic acid is used in the manufacture of nylon-66. Esters of benzoic acid are used in perfumery. Sodium benzoate is used as a food preservative. Higher fatty acids are used for the manufacture of soaps and detergents.

By going through these CBSE Class 12 Chemistry Notes Chapter 11 Alcohols, Phenols and Ethers, students can recall all the concepts quickly.

Alcohols, Phenols and Ethers Notes Class 12 Chemistry Chapter 11

→ Alcohol contains one or more hydroxyl (OH) group (s) directly attached to a carbon atom (s).

→ A Phehol contains one or more hydroxyl group (OH) attached to a carbon atom (s) of the benzene ring.

→ An ether contains an alkoxy/aryloxy group (R-O/Ar-O) in place of the H atom of a hydrocarbon.

Classification:

• C₂H₅OH: Monhydric alcohol
• CH₂OH-CH₂OH: Dihydric alcohol
• HOH₂C-CHOH-CH₂OH: Trihydric alcohol
• 
  

1. Compounds containing C_sp³ -OH bond: In this class of alcohols, the -OH group is attached to an sp³ hybridized carbon atom of an alkyl group. They are further classified as follows:

→ Primary, secondary and tertiary alcohols: Here -the OH group is attached to a primary, secondary and tertiary carbon atom, respectively as shown below:

→ Allylic alcohols: In these alcohols, -OH group is attached to an sp³ hybridized carbon next to the carbon-carbon double bond, i.e., to an allylic carbon, e.g.,

→ Benzylic alcohols: In these alcohols, the -OH group is attached to an sp³-hybridized carbon next to an aromatic ring. For example

Allylic and benzylic alcohols may be primary, secondary or tertiary.

2. Compounds containing C_sp²-OH bond: These alcohols contain -OH group bonded to a carbon-carbon double bond.
Vinylic alcohols CH₂ = CH-OH

Ethers (A) Simple ethers/Symmetrical ethers (ROR)-
C₂H₅-O-C₂H₅ CH₃-O-CH₃.

(B) Mixed or Unsymmetrical ether (ROR’): R ≠ R’
C₂H-O-CH₃ C₆H₅-O-C₂H₅

Nomenclature:
(a) Alcohols: Common and I.U.P.A.C. names given below:

Table 11.1: Common and IUPAC Names of Some Alcohols:

Cyclic alcohols are named using the prefix cyclo and considering the-OH group attached to C-1.

(b) Phenols: The simplest hydroxy derivative of benzene is phenol, it is its common name and also an accepted IUPAC name. As the structure of phenol involves a benzene ring, in its substituted compounds the terms ortho (1, 2- disubstituted), meta (1, 3-disubstituted) and para (1, 4-disubstituted) are often used in the common names.

Dihydroxy derivatives of benzene are known as 1, 2-, 1, 3- and 1, 4-benzenediol.

(c) Ethers: Common names of others are derived from the names of alkyl/aryl groups written as separate words in alphabetical order and adding the word ‘ether’ at the end. For example, CH₃OC₂H₅ is ethyl methyl ether. If both the alkyl groups are the same, the prefix ‘di’ is added before the alkyl group. For example, C₂H₅OC₂H₅ is diethyl ether.

According to the IUPAC system of nomenclature, others are regarded as hydrocarbon derivatives in which a hydrogen atom is replaced by an -OR or -OAr group, where R and Ar represent alkyl and aryl groups, respectively. The larger (R) group being chosen as the parent hydrocarbon. The names of a few others are given as examples in the Table below:

Table 11.2: Common and IUPAC Names of Some Ethers:

Structures of Functional Groups-In alcohols, the oxygen of the -OH group is attached by a sigma (a) bond formed by the overlap of an sp³ hybridized orbital of carbon with an sp³ hybridized orbital of oxygen.

The figure below depicts the structural aspects of the methanol, phenol and methoxymethane.

Structures of methanol, phenol and methoxymethane.

The bond angle in alcohols is slightly less than the tetrahedral angle (109°28′) which is due to repulsion between the unshared electron pairs of oxygen.

Alcohols and Phenols:
Preparation of Alcohols:
1. From alkenes:
1. By acid-catalysed hydration

Mechanism:
→ Step I: Protonation of alkene to form carbocation by the electrophilic attack of H₃O⁺ ion,
H₂O + H⁺ → H₃O⁺

→ Step II: Nucleophilic attack of water on carbocation

→ Step III: Deprotonation to form an alcohol

3. By hydroboration-oxidation: Borane (BH₃) is an electrophile since it is electron-deficient. Addition product formed is oxidized to alcohols by H₂O₃  and aq. NaOH.

Addition of water proceeds against the Markovnikor rule.

2. From Carbonyl Compounds:
1. Reduction of aldehydes and Ketones:
(a) Hydrogen gets added in the presence of a catalyst (catalytical hydrogenation) like Pd, Pt, Ni (all finely. divided)
(b) By treating carbonyl compounds with sodium borohydride or lithium aluminium hydride (Li A1H4).

Aldehydes give primary alcohols whereas ketones yield secondary alcohols.

2. Reduction of carboxylic acids and esters:

3. From Gngnard reagents:


Preparation of Phenols:
1. From haloarenes

2. From benzene Suiphonic acid

3. From diazonium salts

4. From Cumene

→ Physical Properties of Alcohols and Phenols: The properties of alcohols and phenols are chiefly due to the hydroxyl group. The nature of alkyl and aryl groups simply modify their properties.

→ Boiling Points: The boiling points of alcohols and phenols increase With the increase in the number of carbon atoms (increase in van der Waals forces). In alcohols, the boiling points decrease with the increase of branching in the carbon chain (because of a decrease in van der Waals forces due to a decrease in surface area).

The -OH group is alcohols and phenols is involved in intermolecular hydrogen bonding as shown below:

The high boiling points of alcohols are mainly due to the presence of intermolecular hydrogen bonding in them which is lacking in ethers and hydrocarbons.

→ Solubility: Solubility of alcohols and phenols in water is due to their ability to form hydrogen bonds with water molecules. The solubility decreases with an increase in the size of alkyl/aryl (hydrophobic) groups. Lower alcohols are miscible with water in all proportions.

→ Chemical Reactions: Alcohols react both with nucleophiles and electrophiles. The bond between O-H is broken when alcohols react as nucleophiles.
1.

2. The bond between C-O is broken when they react as electrophiles. Protonated alcohols react in this manner.

Based on the cleavage of O-H and C-O bonds, the reactions of alcohols and phenols may be divided into two groups.
(a) Reactions involving cleavage of O-H bond:
1. Acidity of alcohols and phenols

→ Reactions with metals:
2R-O-H + 2 Na → 2 RONa + H₂

→ Reaction with Aq. NaOH: Phenols react with aq. NaOH

→ Alcohols and Phenols are Bronsted acids i.e., they can donate a proton to a stronger base [: B]

→ The acidity of alcohols:

Due to the electron-releasing group (-CH₃, -C₂H₅) electron density on oxygen increases and the polarity of the O-H bond decreases. This decreases the acid strength.

Alcohols are, however, weaker acids than water.

Water is a better proton donor (i.e., stronger acid) than alcohol. Alkoxides (e.g. sodium ethoxide) is a better proton-acceptor than hydroxide ion, which suggests that alkoxides are stronger bases. (Sodium ethoxide is a stronger base than sod. hydroxide).

Alcohols act as Bronsted bases as well. It is due to the presence of unshared electron pairs on oxygen, which makes them proton acceptors.

→ The acidity of Phenols: The reactions of phenols with metals like Na, Al and NaOH indicate their acidic nature. The -OH group in phenols is directly attached to the sp2 hybridized carbon of benzene ring which acts as an electron-withdrawing group.

The reaction of phenol with aqueous sodium hydroxide indicates that phenols are stronger acids than alcohols and water. Let us examine how a compound in which the hydroxyl group attached to an aromatic ring is more acidic than the one in which the hydroxy] group is attached to an alkyl group.

The ionisation of alcohol and a phenol takes place as follows:

Due to the higher electronegativity of sp² hybridised carbon of phenol to which -OH is attached, electron density decreases on oxygen. This increases the polarity of the O-H bond and results in an increase in ionisation of phenols than that of alcohols. Now let us examine the stabilities of alkoxide and phenoxide ions. In alkoxide ion, the negative charge is localised on oxygen while in phenoxide ions, the charge is delocalised.

The delocalisation of negative charge (structures I-V) makes phenoxide ion more stable and favours the ionisation of phenol. Although there is also charge delocalisation in phenol its resonance structures have charge separation due to which the phenol molecule is less stable than phenoxide ion.

In substituted phenols, the presence of electron-withdrawing groups such as the nitro group enhances the acidic strength of phenol. This effect is more pronounced when such a group is present at ortho and para positions. It is due to the effective delocalisation of negative charge in phenoxide ion.

On the other hand, electron releasing groups, such as alkyl groups, in general, do not favour the formation of phenoxide ion resulting in a decrease in acid strength. Cresols, for example, are less acidic than phenol.

Table 11.3: pKa Values of Some Phenols and Ethanol:

Compound	Formula	pKa
o-Nitrophenol	0-O2N-C6H4-OH	7.2
m-Nitrophenol	m-O2N-C6H4-0H	8.3
p-Nitrophenol	p-O2N-C6H4-0H	7.1
Phenol	C6H5-OH	10.0
o-Cresol	O-CH3-C6H4-OH	10.2
m-Cresol	m-CH3C6H4-OH	10.1
p-Cresol	p-CH3-C6H4-OH	10.2
Ethanol	C2H5OH	15.9

From the above data, you will note that phenol is a million times more acidic than ethanol.

2. Esterification: Alcohols and phenols react with carboxylic acids, acid chlorides and acid anhydrides to form esters.

The introduction of acetyl (CH₃ CO^–) group in alcohols or phenols is known as acetylation. Acetylation of salicylic acid produces aspirin which possesses analgesic, anti-inflammatory and antipyretic properties.

(b) Reactions involving cleavage of C-O bond in alcohols.
1. Reaction with HX:
ROH + HX → R-X + H₂O.

Luca’s Test distinguishes the three classes of alcohols (1°, 2° and 3°) on reaction with cone. HCl and ZnCl₂ [Luca’s reagent], 3° alcohols produce turbidity immediately with it. 2° alcohols do it after some time. 1° alcohol does not produce turbidity at room temperatures.

2. Reactions with phosphorus trihalides:
3R-OH + PCl₃ → 3R-Cl + H₃P0₃

3. Reaction with a protic acid: e.g., cone. H₃P0₄ or H₂SO₄ causes dehydration of alcohols producing alkenes.

Thus the relative ease of dehydration of alcohols follows the order Tertiary > Secondary > Primary

Mechanism of dehydration of alcohols:
→ Step I: Formation of protonated alcohol

→ Step II: Formation of the carbocation. It is the slow step and hence the rate-determining step.

→ Step III: Formation of ethene by loss of a proton.

4. Oxidation:

(PCC is pyridinium chlorochromate-a complex of chromium oxide with pyridine and HCl)

Tertiary alcohols do not undergo oxidation reactions.

5. Dehydrogenation with red hot copper: I° and 2c alcohol form aldehydes and ketones respectively. 3° alcohols undergo dehydration.

(c) Reactions of Phenols:
Following reactions are shown by phenols only.
1. Electrophilic aromatic substitution: The —OH group fused in the ring activates the ring towards electrophilic substitution directing the incoming group at ortho and para positions.
1. Nitration:

The o- and p-isomers can be separated by steam distillation, o-, Nitrophenol is steam volatile due to intramolecular H-bonding while p-nitrophenol (Higher B.Pt) is less volatile due to intermolecular H- bonding which causes the association molecules.

2. Halogenation:

Reaction with Bromine water is used as a test of phenol.

2. Koibe’s reaction:

3. Reimer-Tiemann reaction: Salicylaldehyde is formed.

4. Reaction with zinc dust: Benzene is formed.

5. Oxidation:

Some Commercially Important Alcohols:
1. Methanol (CH₃OH): It is also known as wood-spirit.

Properties: It is a colourless liquid. It boils at 337 K. It is highly poisonous in nature. Ingestion of even small quantities causes blindness and large quantities cause even death.

Uses:

• It is used as a solvent in paints, varnishes etc.
• It is chiefly used for making formaldehyde.

2. Ethanol (C₂H₅OH):
Commercial Preparation: By fermentation of molasses

Properties:

1. It is a colourless liquid with B.Pt 351 K.
2. Combined with CuSO₄ and pyridine, it is termed as denatured spirit

Uses:

• It is an excellent solvent.
• In the laboratory and hospitals for sterilisation of surgical instruments.

Ethers:
Preparation of Ethers:
1. By dehydration of alcohols:

It is a nucleophilic bimolecular reaction (SN2)

Step:


2. Willamson synthesis of ethers: It is an important laboratory method for the preparation of symmetrical and unsymmetrical ethers.

It involves the S_N² attack of an alkoxide ion on primary alkyl halide.

Physical Properties of ethers:

1. The C-O bonds in ethers are polar and thus ethers have a net dipole moment.
2. Their b. pts are comparable to those of alkanes with comparable molecular masses but lower than those of alcohols. It is due to the lack of H-bonding in ethers.
3. The miscibility of ethers with water resembles those of alcohols of the same molar mass. It is due to the fact that-ethers like alcohols can form H-bonds with water.
   

Chemical Reactions: Ethers are the least reactive of the functional groups.
1. Cleavage of C-O bond in ethers-lire the cleavage of C-O bond in ethers takes place under drastic conditions with an excess of hydrogen halides.

The order of reactivity is HI > HBr > HCl.

Mechanism:
→ Step I:

→ Step II:

With HI in excess and at a high temp., ethanol reacts with another molecule of HI and is converted to ethyl iodide.

→ Step III:

However, when one of the alkyl groups is tertiary, the halide formed is a tertiary halide.

It is because, in step-2 of the reaction, the departure of leaving group (HO- CH₃) creates a more stable carbocation [(CH₃)₃ C⁺] and the reaction follows the S_N¹ mechanism.

2. Electrophilic substitution: The alkoxy group (-OR) is ortho and para directing and activates the aromatic ring towards electrophilic substitution.

1. Halogenation: Phenylalkyl ethers undergo usual halogenation in the benzene ring, e.g., anisole undergoes bromination with bromine in ethanoic acid even in absence of iron (III) bromide catalyst. It is due to the activation of the benzene ring by the methoxy group. Para isomer is obtained in 90% yield.

2. Friedel Crafts reaction: Anisole undergoes Friedel Crafts reaction, i.e., the alkyl and acyl groups are introduced at ortho and para positions by reaction with an alkyl halide and acyl halide in the presence of anhydrous aluminium chloride (a Lewis acid) as a catalyst.


3. Nitration: Anisole reacts with a mixture of the cone. H₂SO₄ and HNO₃ to yield a mixture of ortho and para nitro anisole.

By going through these CBSE Class 12 Chemistry Notes Chapter 10 Haloalkanes and Haloarenes, students can recall all the concepts quickly.

Haloalkanes and Haloarenes Notes Class 12 Chemistry Chapter 10

In Haloalkanes X is attached to sp³ hybridized carbon atom, whereas it is attached to sp² hybridized carbon atom in the aryl group.

Classification:

(a) Alkyl halides or Haloalkanes (R-X) [sp³ C-X Bond]
General Formula: C_n H_2n-1X

(b) Allylic halides: Here halogen atom is bonded to an sp³-hybridized carbon atom next to C = C, i.e., to an allylic carbon.

(c) Benzylic halides: Halogen is bonded to sp³ carbon next to the aromatic ring.

Compounds containing sp² C-X Bond:



IUPAC name:

Common and IUPAC names of some halides:

Structure	Common name	IUPAC name
CH2Cl2	Methylene chloride	Dichloromethane
CHCl3	Chloroform	Trichloromethane
CHBr3	Bromoform	Tribromomethane
CCI4	Carbon tetrachloride	Tetrachl or methane
CH3CH2CH2F	n-Propyl fluoride	1-Fluoropropane

Nature of C-X bond: Due to the difference in electronegativity of C and X, the C-X bond is polarised; carbon bears a partial positive charge whereas the halogen atom bears a partial negative charge.

Carbon-halogen bond length increases from C-F to C-I as the size of the halogen atom increases.

Methods of Preparation:
1. From alcohols

The order of reactivity of alcohols with a given haloacid is 3° > 2° > 1°.

2. From hydrocarbons:
1. By Free radical halogenation: It gives a complex mixture of isomeric mono and polyhaloalkanes which is difficult to separate.

2. By electrophilic Substitution: Aryl chlorides and bromides can be easily prepared by electrophilic substitution of arenes with chlorine and bromine respectively in the presence of Lewis acid catalysts like iron or iron (III) chloride.

The ortho and para isomers can be easily separated due to large differences in their melting points.

3. Sandmayer’s reaction:

4. From alkenes: (a) Addition of hydrogen halides

Addition to unsymmetric alkenes is as per Markovnikov’s Rule

(b) Addition of Halogens

5. Halogen Exchange: Finkelstein Reaction
R-X + Nal → R-I + NaX
X = Cl, Br

Swartz Reaction: This method is used to prepare alkyl fluorides by heating an alkyl chloride/bromide in the presence of AgF/Hg₂F₂.
H₃C-Br + AgF → H₃C-F + AgBr

Physical Properties:

1. Alkyl halides are colorless when pure. However, bromides and iodides develop color when exposed to light.
2. Melting & b.Pts: Lower members are gases at room temperature. Higher members are liquids or solids.

Due to the polar character of the C-X bond and higher molecular mass as compared to the parent hydrocarbon, the intermolecular forces of attraction (dipole-dipole and van der Waals) are stronger in halogen derivatives. That is why boiling points of chlorides, bromides, and iodides are considerably higher than those of the hydrocarbons of comparable molecular mass.

For the same alkyl group, the boiling points increase from RF to RI in the order RF < RCl < RBr < RI.
For isomeric haloalkenes, the b.pts decrease with an increase in branching (lesser the surface area)

B. pts of isomeric di-halogens are very nearly the same. However, para isomers are higher melting as compared to their ortho and meta isomers. It is due to the symmetry of para isomers that fits in the crystal lattice better as compared to ortho and meta isomers.

3. Density: Bromo, iodo, and poly-chloro derivatives of hydrocarbons are heavier than water. The density increases with an increase in the number of carbon atoms, halogen atoms, and atomic mass of halogens.

4. Solubility: The haloalkanes are only very slightly soluble in water. However, they tend to dissolve in organic solvents.

Chemical Reactions:
A. Reactions of haloalkanes:

1. Nucleophilic substitution reactions (S_N)
2. Elimination reactions
3. Reactions with metals.

1. Nucleophilic Substitution reactions: A nucleophile (Nu^– 🙂 reacts with haloalkane (the substrate) which has a polar C-X bond.



Such reactions in which a stronger nucleophile displaces a weaker nucleophile are called Nucleophilic Substitution (S_N) reactions and the halide ion which departs with its bonding pair of electrons is called the leaving group. The better the leaving group, the more facile is the nucleophile substitution reaction. It follows the order
I^– > Br > Cl^– > F^–
∴ The order of reactivity of haloalkanes follows the sequence Iodoalkanes > Bromoalkanes > chloroalkanes > fluoroalkanes.

Types of Nucleophilic Substitution reactions:

1. S_N² [Substitution, nucleophilic, bimolecular]
2. S_N¹ [Substitution, nucleophilic, unimolecular]

1. Substitution nucleophilic bimolecular (S_N²): The reaction between CH₃Cl and hydroxide ion to yield methanol and chloride ion follows second-order kinetics, i.e., the rate depends upon the concentration of both reactants.

rate of reaction ∝ [Base] [R-X]

Since the rate of the reaction depends upon the concentration of both the reactants, it is a bimolecular nucleophilic displacement reaction.

There occurs a complete stereochemical inversion of the configuration. The order of reactivity of the alkyl halides is Primary halide > Secondary halide > Tertiary halide.

In the S_N² reaction, the attack of the nucleophile (OH^– above) occurs from the backside, and the halide ion leaves from the front side. This inversion of configuration is called Walden Inversion. As far as the ease of departure of halide ion is concerned, the order of reactivity is RI > RBr > RCl > RF.

2. Substitution, nucleophilic, unimolecular (S_N¹) – S_N¹ reactions are generally carried out in polar protic solvents (like water, alcohol, acetic acid, etc). The reaction between tert-butyl bromide and hydroxide ion yields tert-butyl alcohols and follows first-order kinetics.

This reaction is independent of the concentration of the base. The rate law suggests the reaction proceeds in two steps.

This step is slow and hence is the rate-detaining step.

This step, being fast, does not affect the rate of reaction. If the alkyl halide is optically active, then the product is a racemic mixture.

Allylic and benzylic halides show high reactivity towards SN1 reaction

The carbocation gets stabilized through resonance as shown below:

For a given alkyl group, the reactivity of the halide, R-X, follows the same order in both mechanisms.
R-I > R-Br > R-Cl > > R-F.

Stereochemical aspects of nucleophilic substitution reactions: An S_N² reaction proceeds with complete stereochemical inversion of configuration while an S_N¹ reaction proceeds with racemization.

Optical Activity: Certain compounds exhibit the property of rotating the plane polarised light when it passed through their solutions. Such compounds are called Optically Active compounds arid this phenomenon is called Optical Activity.

If the compound rotates the plane-polarised light to the right, i.e., in a clockwise direction, it is called dextro-rotatory or the d-form and is indicated by placing a positive (+) sign before the degree of rotation. If the light is rotated towards the left (anticlockwise), the compound is said to be laevorotatory or the /-form and a negative sign (-) is placed before the degree of rotation. Such (+) and (-) isomers of a compound are called Optical Isomers and this phenomenon is termed Optical Isomerism.

All the physical properties of compounds showing optical activity are the same like refractive index solubility, density, m.pts, b.pts, etc. Even the extent of rotation is the same. They differ from each other only in the direction of rotation.
If all the substituents attached to the C atom are different, such a carbon atom is called asymmetric carbon or stereocentre.

The resulting molecule would lack symmetry and is referred to as asymmetric or ‘dissymmetric molecule. This asymmetry of the molecule is responsible for the optical activity in such organic compounds. Such molecules are non-superimposable on their mirror image (as the left hand is non-superimposable on the right hand) and are said to be Chiral. This property of non-super imposibility of the mirror image on the object is called Chirality. The objects which are superimposable on their mirror image are called Achiral.

Butan-2-ol has 4 different groups attached to the tetrahedral carbon atom and is Chiral. The mirror image of Butan-2-ol non-superimposable onbutan-2-ol

Other chiral molecules are Bromochloroiodimethane (BrCl CHI), 2-chlorobutanol, 2, 3-dihydroxypropanal (OHC- CHOH-CH₂OH), lactic acid (CH₃CH(OH)COOH). The stereoisomers related to each other as non-superimposable are also called Enatiomers and the concept Enantiomerism.

A mixture containing two enantiomers in equal proportions will have zero optical rotation, as the rotation due to one isomer will be canceled by the rotation due to the other isomer. Such a mixture is called Racemic Mixture or Racemic Modification. It is represented by prefixing dl or (±) before the name, e.g., (±) butan-2-ol. The process of conversion of enantiomer into a racemic mixture is known as Racemisation.

Retention: Retention of configuration is the preservation of the integrity of the spatial arrangement of bonds to an asymmetric center during a chemical reaction or transformation. It is also the configurational correlation when a chemical species XCabc is converted into the chemical species YCabc having the same relative configuration.

Inversion, retention, and racemization: These are three possibilities for a reaction to occur at an asymmetric carbon atom.

Consider the replacement of a group X and Y in the following reaction:

If (A) is the only compound obtained, the process is called retention of configuration.
If (B) is the only compound obtained, the process is called inversion of configuration.
If a 50: 50 mixture of (A) and (B) is obtained, the process is called racemization and the product is optically inactive.

Thus during an S_N² reaction involving an optically active alkyl halide, the reactant undergoes inversion of configuration.

In the case of optically active alkyl halides, S_N¹ reactions are accompanied by racemization. Consider hydrolysis of optically active 2-romobutane, which results in the formation of(±) butan-2-oI.

Step I:


2. Elimination Reactions: When a haloalkane with a β-hydrogen atom is heated with an alcoholic solution of potassium hydroxide, there is an elimination of hydrogen atom from β-carbon and a halogen atom from the a-carbon atom.

An alkene is formed as a result. Since the β-hydrogen atom is involved in elimination, it is often called β-elimination.

If there is the possibility of the formation of more than one alkene due to the availability of more than one β-hydrogen atom, usually one alkene is formed as the major product. These form part of a pattern first observed by Russian Chemist Alexander Zaitsev (also pronounced as Saytzeff) who in 1875 formulated a rule which can be summarised as in dehydrohalogenation reactions, the preferred product is that alkene which has the greater number of alkyl groups attached to the doubly bonded carbon atoms.” Thus, 2-bromopentane gives pent-3-ene as the major product.

An alkyl halide with β-hydrogen atoms when reacted with a base or a nucleophile has two competing routes: substitution (S_N¹ and S_N²) and elimination. The route to be taken up depends upon the nature of alkyl halide, strength and size of base/nucleophile, and reaction conditions.

Thus, a bulky nucleophile abstracts a proton rather than approaches a tetravalent C atom (steric hindrance). Similarly, a primary alkyl halide will prefer an S_N² reaction, a secondary halide S_N² or elimination depending upon the strength of base/nucleophile, and a tertiary halide: S_N¹ or elimination depending upon the stability of carbocation or the more substituted alkene.

3. Reaction With Metals: Most organic chlorides bromides and iodides react with certain metals to give compounds containing carbon-metal bonds. Such compounds are known as organometallic compounds.

In the Grignard reagent, the C-Mg bond is covalent but highly polar, the MgX bond is essentially ionic.

The Grignard reagent is highly reactive. Even H₂O reacts with it.

Wurtz Reaction: Alkyl halides react with sodium in dry ether to give hydrocarbons containing double the number of carbon atoms present in the reaction. This reaction is known as the Wurtz reaction.

Reactions of Haloarenes:
1. Nucleophilic Substitution: Alkyl halides are extremely dull/ loss reactive towards S_N reactions due to the following reasons.
1. Resonance effect: In haloarenes, the electron pairs on halogen atom are in conjugation with π-electrons of the ring as given below:

C-Cl bond a partial double bond character due to one. As a result difficult to break the C-X bond and therefore less reactive towards S_N¹
2. Different hybridization of Carbon atom in C-X bond: In haloalkanes, the C is sp³ hybridized while in haloarenes, the carbon attached to has is sp² hybridized.

The sp² hybridized carbon with \(\frac{1}{3}\) s-character is more electronegative and can hold the electron pair of C-X bond more tightly than sp3-hybridized carbon in haloalkane ‘With s-character. Thus

The C-Cl bond length in haloalkane is 177 pm while in haloarene it is 169 pm. Since it is difficult to break a shorter bond than a longer bond, therefore, haloarenes are less reactive than haloalkanes towards SN reactions.

3. Instability of phenyl cation: In the case of haloarenes, the phenyl cation formed as a result of self-ionization will not be stabilized by resonance and, therefore, the S_N¹ mechanism is ruled out.

4. Because of the possible repulsion, it is less likely for the electron-rich nucleophile to approach electron-rich arenes.
Replacement by hydroxyl group:

The presence of an electron-withdrawing (-NO₂) at ortho and para position increases the reactivity of haloarenes.
The effect is pronounced when the (-NO₂) group is introduced at ortho and para positions. However, no effect on the reactivity of haloarenes is observed by the presence of an electron-withdrawing group at meta-position. Mechanism of the reaction is as depicted

The effect is pronounced when (—NO₂ the) group is introduced at ortho and para positions. However, no effect on the reactivity of haloarenes is observed by the presence of an electron-withdrawing group at the meta position. The mechanism of the reaction is as depicted:



The presence of a nitro group at ortho- and para- position withdraws the electron density from the benzene ring and thus facilitates the attack of the nucleophile on haloarene.

The carbanion thus formed is stabilized through resonance. The negative charge appeared at ortho- and para- positions with respect to the halogen substituent is stabilized by -NO₂ group while in the case of mcfa-nitrobenzene, none of the resonating structures bear the negative charge on carbon atom bearing the -NO₂ group.

Therefore, the presence of a nitro group of meta-position does not stabilize the negative charge and no effect on reactivity is observed by the presence of the -NO₂ group of meta-position.

2. Electrophilic Substitution Reactions: Haloarenes undergo the usual electrophilic reactions of the benzene ring such as halogenation, nitration, sulphonation, and Friedel-Crafts reactions. Halogen atom besides being slightly deactivating is o, p-directing, therefore, further 1 substitution occurs at ortho- and para-positions with respect to the halogen atom.

The o, p-directing; therefore, further substitution occurs at ortho- and para-positions with respect to the halogen atom. The o, p-directing influence of halogen atom can be easily understood if we consider the resonating structures of halobenzene as shown:

Due to resonance, the electron density increases more at ortho- and para-positions than at mcia-positions. Further, the halogen atom because of its-I effect has some tendency to withdraw electrons from the benzene ring. As a result, the ring gets somewhat deactivated as compared to benzene and hence the electrophilic substitution reactions in haloarenes occur slowly and require more drastic conditions as compared to those in benzene.
1. Halogenation:

2. Nitration:

3. Sulphonation:

4. Friedel-Crafts reaction:


3. Reaction with metal:
Wurtz-Fitting Reaction: It is between an alkyl and aryl halide.

→ Polyhalogen Compounds: Carbon compounds containing more than one halogen atom are usually referred to as polyhalogen compounds.

→ Dichloromethane (Methylene chloride CH₂Cl₃: It is widely used as a solvent in paint remover, as a propellant in aerosols, and as a process solvent in the manufacturing of drugs. It harms the human central nervous system.

→ Trichloromethane (Chloroform CHCl₃): It is used as a solvent for fats, alkaloids, iodine, waxes, rubbers, plastics, etc. It was used as a general anesthetic in surgery.

Chloroform is slowly oxidized to carbonyl chloride (phosgene) by air in the presence of light. It is extremely poisonous in nature.

It is therefore stored in colored bottles to cut off light and in well- stoppered fully filled bottles to cut off air.

→ Triiodomethane (Iodoform CHI3₃): It was used as an antiseptic, but the antiseptic properties are due to the liberation of free iodine and not due to the iodoform itself.

→ Tetrachloromethane (Carbon Tetrachloride CCl₄): It is used for the synthesis of chlorofluorocarbons and as a solvent. There is some evidence that exposure to CCl₄ causes liver cancer in humans. It causes dizziness, nausea, and vomiting which can cause permanent damage to nerve cells. The chemical may irritate the eyes on contact. It depletes the ozone layer when released into the air.

→ Freons: Chlorofluorocarbon compounds of methane and ethane are collectively known as freons. They are extremely stable, unreactive, non-toxic, non-corrosive, and easily liquefiable gases, Freon 12 (CCl₂F₂) is one of the most common freons in industrial use. Most freon, even that used in refrigeration, eventually makes its way into the atmosphere where it diffuses into the stratosphere where it is able to imitate radical change reactions that can upset the natural ozone balance.

→ p, p’-Dichlorodiphenyltrichloroethane (DDT): The use of DDT was effectively used against the mosquito that spreads malaria and lice that carry typhus. Many species of insects developed resistance to DDT, and DDT was also discovered to have a high degree of toxicity towards fish. The chemical stability of DDT and its fat solubility compounded the problem. The use of DDT was banned in the USA in 1973, although it is still in use in some parts of the world. Its chemical formula is

Many out of these polyhalogen compounds cannot be easily decomposed and cause of depletion of the ozone layer and are proving environmental hazards.

By going through these CBSE Class 12 Chemistry Notes Chapter 8 The d-and f-Block Elements, students can recall all the concepts quickly.

The d-and f-Block Elements Notes Class 12 Chemistry Chapter 8

Transition Elements: d-Block. elements are called transition elements. They are placed in between the s-block on their left and the p-block on their right. They are called transition elements because they have properties intermediate between those of s- and p-block elements and represent a change from the most electropositive s-block elements to the least electropositive elements.

This element has partly filled (n -1) d-subshell in their atomic or ionic state. They are called d-block elements since in them 3d, 4d, 5d and 6d subshells are incomplete and the last electron enters (n -1) d subshell, i.e., penultimate (last but one) shell.

General Electronic Configuration:
The outermost shell: Their general electronic configuration is (n – 1) d^1-10 ns^0-2 where n is the outermost shell. Filling of 3d, 4d, 5d and 6d subshells leads to first transition series [4th period], 2nd transition series [5th period], 3rd transition series [6th period] and 4th incomplete transition series [7th period] respectively.

1st Transition Series or 3d series: Starts from Sc (At. No. 21) and ends at Zinc (Zn; atomic no. = 30).

Their configuration is 3d^1-10 ns^1-2.

• Scandium (Sc) = 21 = [Ar]¹⁸ 3d¹ 4s².
• Titanium (Ti) = 22 = [Ar]¹⁸ 3d² 4s²
• Vanadium (V). = 23 = [Ar]¹⁸ 3d³ 4s²
• Chromium (Cr) = 24 = [ Ar]¹⁸ 3d⁵ 4s¹
• Manganese (Mn) = 25 = [Ar]¹⁸ 3d⁶ 4s²
• Iron (Fe) = 26 = [Ar]¹⁸ 3d⁶ 4s²
• Cobalt (Co) = 27 = [Ar]¹⁸ 3d⁷ 4s²
• Nickel (Ni) = 28 = [Ar]¹⁸ 3d⁸ 4s²
• Copper (Cu) = 29 = [Ar]¹⁸ 3d¹⁰ 4s¹
• Zinc (Zn) == 30 = [Ar]¹⁸ 3d¹⁰ 4s²

It may be noted that the electronic configuration of chromium and copper given above are exceptional. They have a single electron in the 4s-orbital. It is due to the extra-stability of 3d⁵ and 3d¹⁰ (half-filled and completely filled) orbitals. They consist up of 10 elements.

2nd Transition Series or 4d Series: Starts from the element Ytterium (Y) [Z = 39] and ends of cadmium (Cd; Z = 46). Their electronic configuration is 4d^0-10 5s^1-2. Palladium (Pd; Z = 46) has exceptional configuration of 4d¹⁰ 5s°. They consists up of ten elements.

Third Transition Series or 5d series: Corresponds to filling of 5d sub-level. They consist up of ten elements La (Z = 57), H/(Z = 72), Ta, W, Re, Os, Ir, Pt, Au and Hg(Z = 80).

The electronic configuration of 3rd transition elements in the inner and valence shell is given below:

Fourth Transition Series or 6d Series: Corresponds to tire filling up of 6d Sub level and starts with Actinium (Ac; Z = 89), Rf (earlier Ku, Z = 104), Db, Sg, Bh, Hs, Mt, Ds, Rg and ends at Uub (Z = 112).

The fundamental difference in the electronic configuration of Representative Elements and Transition Elements: In the representative elements (s- and p-block elements), the valence electrons are present only in the outermost shell while in the transition elements, the valence electrons are present in the outermost shell as well as d- orbitals of the penultimate shell.

Zinc (Zn), Cadmium (Cd) and mercury (Hg) are misfits according to the definition of transition elements as they have 3d¹⁰, 4d¹⁰, 5d¹⁰ (completely filled inner d-orbitals) in the ground state of their atoms or in one of the common oxidation states (dipositive ions: Zn²⁺, Cd²⁺, Hg²⁺) respectively. They do not show properties of transition elements to any appreciable extent, except for their ability to form complexes.

The d-block of the periodic table contains the elements of Groups 3-12 in which the d-orbitals progressively filled in each of the four long periods.

General Properties of Transition Elements:
(a) Metallic character: All the elements of d block are metals. This is due to the fact that they have low ionisation energy values and have only one dr two s-electrons in the outermost shell. The transition elements exhibit good mechanical properties i.e., they are hard, malleable and ductile. Except for Zn, Cd and Hg, they have high m. and b.pts. They have thermal and electrical conductivity and metallic lustre.

(b) Variable oxidation states: The transition elements exhibit a variety of oxidation states in their compounds. This is due to the fact that (n -1) d orbitals are of comparable energy to ns orbitals and therefore some or all of the (n -1) d electrons can be used along with ns electrons in compounds formation. Some common oxidation states exhibited by elements of the first transition series are listed below:

(The value in parentheses are less common oxidation states.)

(c) Coloured compounds: Most of the compounds of transition elements are coloured. The d orbitals in transition metal compounds are not of equal energy. In transition elements, the d-orbital are partly filled and electron may be promoted from one d level to another d level by absorbing visible light. Consequently, the compound has the colour complement to the absorbed light. For example, Cu₂ absorbs red light and transmitted light contains an excess of the other colours of the spectrum and appears to be blue.

The ions Zn²⁺, Cu⁺, Ti⁴⁺, Sc³⁺ are white because they have either completely empty or completely filled d-subshell.

(d) Magnetic properties: Most of the transition metals and their compounds are paramagnetic i.e., they are attracted by the magnetic field. This, the property is due to the presence of unpaired electrons in the d orbitals of their atoms. The elements like iron, nickel and cobalt have appreciable paramagnetism and are called ferromagnetic substances. The magnetic moment of compounds of d-block elements are determined by spin only values which may be related to the number of impaired electrons (n) as

Spin magnetic moment (μ_s) = \(\sqrt{n(n+2)}\)

The agreement between the observed and calculated values is quite good for most of the elements. However, for the latter half of the series, the contribution from orbital magnetic moments is observed.

(e) Complex formation: The transition elements have a great tendency to form complexes.

This is because of:

• their small cation size
• high effective nuclear charge and
• presence of empty d-orbitals of appropriate energy for bonding to the ligands.

It has been noticed that

1. In each transition series, the stability of complexes increases with an increasing atomic number of the element and in a particular oxidation state with decreasing size of its atoms.
2. In the case of metal atoms that show more than one oxidation state, the complexes of greater stability are formed with the metal ions of the highest charge.

(f) Catalytic properties: Many transition metals and their compounds show catalytic properties e.g., Fe, Pt, V₂O₅, Ni etc. This property may be due to either the use of the d-orbitals or from the formation of transient intermediate compounds to absorb and activate the reacting substances.

(g) Alloy Formation: Transition metals mix freely with each other in the molten state and on cooling a solution of different transition metals to form alloys. For example, chromium dissolves in nickel to form Cr-Ni alloy; manganese dissolves in iron to form manganese steel. They take part in alloy formation because their atoms can exchange lattice sites of each other. Such alloys are formed by atoms with metallic radii that are within about 15 per cent of each other.

(h) Formation of Interstitial Compounds: Interstitial compounds are those which are formed when small atoms like H, C or N are trapped inside the crystal lattices of metals. They are usually non-stoichiometric and are neither typically ionic nor covalent. Many of the transition metals form interstitial compounds particularly with small non-metal atoms such as hydrogen, boron, carbon and nitrogen.

These small atoms enter into the voids sites between the packed atoms of the crystalline metal, e.g., TiC, Mn4N, Fe3M and TiH2 etc. The formulae quoted do not, of course, correspond to any normal oxidation state of the metal and often non-stoichiometric material is obtained with such composition as VH056 and TiH17. Because of the nature of their composition, these compounds are referred to as interstitial compounds.

The principal physical and chemical characteristics of these compounds are as follows:

1. They have high melting points, higher than those of pure metals.
2. They are very hard, some borides approach diamond in hardness.
3. They retain metallic conductivity.
4. They are chemically inert.

Some Important Compounds of Transition Metals:
Oxides: Transition metal oxides are formed by the action of oxygen with transition metals at high temperature. The general formulas of the oxides of transition metals are MO, M₂O₃, M₃O₄, MO₂, M₂O₅ and MO₃.

It has been observed that:

• the oxides in which the metals are in a lower oxidation state are basic in nature.
• the oxides in which the metal is in a higher oxidation state are acidic in nature.
• the oxides in which the metal exhibits intermediate oxidation states are amphoteric.

For example, among the oxides of manganese, MnO is basic, Mn₂O₃, Mn₃O₄ and MnO₂ are amphoteric while Mn₂O₇ is acidic.

Oxides Of 3d Metals:

Chemical Reactivity and E° Values: Transition metals vary widely in their chemical reactivity. Many of them are sufficiently electropositive to dissolve in mineral acids, although a few are ‘noble’: that is, they are unaffected by simple acids.

The metals of the first series with the exception of copper are relatively more reactive and are oxidised by 1M H⁺, though the actual rate at which these metals react with oxidising agents like hydrogen ion (H⁺) is sometimes slow. For example, titanium and vanadium, in practice, are passive to dilute non-oxidising acids at room temperature. The Ee values for M²⁺/M indicate a decreasing tendency to form divalent cations across the series.

This general trend towards less negative Ee values is related to the increase in the sum of the first and second ionisation enthalpies. It is interesting to note that the Ee values for Mn, Ni, and Zn are more negative than expected from the general trend. Whereas the stabilities of half-filled d subshell (d⁵) in Mn²⁺ and completely filled d subshell (d¹⁰) in zinc are related to their E^θ values; for nickel, E^θ value is related to the highest negative enthalpy of hydration.

An examination of the E^θ values for the redox couple M³⁺/M²⁺ shows that Mn³⁺ and CO³⁺ ions are the strongest oxidising agents in aqueous solutions. The ions Ti²⁺, V²⁺ and Cr²⁺ are strong reducing agents and will liberate hydrogen from a dilute acid. e.g.
2 Cr²⁺(aq) + 2 H⁺(aq) → 2 Cr³⁺(aq) + H₂(g)

Trends in Stability of Higher oxidation states as exhibited in the oxides and halides of 3d-Series:
The table below shows the stable halides of the 3d series of transition metals. The highest oxidation numbers are achieved in Ti tetrahalides, VF₅ and CrF₆. The +7 state for Mn is not represented but MnOF₃ is known and beyond Mn, no metal has a trihalide except Fe and COF₃.

The ability of F to stabilize the higher oxidation state is due to either higher lattice energy of the higher compound as in the case of COF₃ or higher bond energy terms for the higher covalent compounds e.g. VF₅ and CrF₆.

Table: Formulae of halides of groups 4 to 12:

Halides: Transition metals react with halogens at high temperature to form transition metal halides. The order of reactivity of halogens is F > Cl > Br > I. The halides of higher oxidation states are the fluorides because metals can be oxidised to higher oxidation states with only fluorine which is most reactive. Bonding in fluorides is mainly ionic. In chlorides, bromides and iodides, the ionic character decrease with the atomic mass of halogen.

The ability of O to stabilize the higher oxidation state is also demonstrated in the oxohalides. Although V^v is represented only by VF₅, the oxohalides VOX₃ are known where X = F, Cl or Br. Another feature of fluorides is their instability in the low oxidation states e.g., VX₂ (X = Cl, Br or I) and the same applies to CuX. On the other hand, all Cu¹¹ halides are known except the iodide, in this case, Cu²⁺ oxidises I^– to I₂:
Cu²⁺ + 2I^–→ CuI(S) + \(\frac{1}{2}\) I₂

However Cul is unstable in solution and undergoes disproportionation.
2Cu⁺ → Cu²⁺ + Cu

The stability of Cu²⁺ (aq) rather than Cu⁺ (aq) is due to the much more negative Ah d H ° of Cu²⁺ (aq) than Cu⁺ (aq), which more than compensates for the 2nd IE of Cu.

The highest oxidation number in the oxides (Table) coincides with the group number and is attained in Sc₂O₃, to Mn₂O₇. Beyond Group 7, no higher oxides of Fe above Fe₂O₃, are known, although ferrates (VI):(FeO₄)^2- are formed in alkaline media they readily decompose to Fe₂O₃ and O₂. Besides the oxides, oxidations stabilise V^v as VO²⁺, Vlv VO²⁺ and TiIV as TiO²⁺.

The ability of O to stabilise these high oxidation states exceeds F in this regard. Thus the highest Mn fluoride is MnF₄ whereas the highest oxide is Mn₂O₇. The ability of oxygen to form multiple bonds to metals explains its superiority. In the covalent oil Mn₂O₇, each Mn is tetrahedrally surrounded by O’s including an Mn-O-Mn bridge. The tetrahedral [MO₄]^n- ions are known for V^v, Cr^vI, Mn^v, Mn^vI and Mn^vII.

Comparison of the First Row Transition Metals Through The d-EIectron Configuration:
In the d° configuration of the simple ions, only Sc³⁺ is known to have this configuration. This configuration then occurs for those metals in which the formal oxidation states equal the total no. of 3d and 4s electron. This is true for Ti (IV), V (V), Cr (VI) and Mn (VII).

→ The d¹ Configuration: Except Vanadium (IV), all others with this configuration are either reducing or undergo disproportionation. For .example, disproportionation occurs for Cr (V) and Mn (VI)
3 CrO₄^3- + 8H⁺ → 2 CrO₄^– + Cr³⁺ + 4 H₂O
3 MnO₄^2- + 4H⁺ → 2 MnO₄^– + MnO₃  + 2H₂0

→ The d² Configuration: This configuration ranges from Ti¹¹ which is very strongly reducing, to FeVI which is very strongly oxidizing. Vanadium (III) is also reducing.

→ The d³ configuration is shown by Chromium (III). It is quite stable and takes part in complex formation.

→ The d⁴ Configuration: There are really no stable species with the configuration Cr (II) is strongly reducing. ,

→ The d⁵ Configuration: The two important species with this configuration are Mn²⁺ and Fe³⁺, the latter may, however, be reduced to Fe²⁺.

→ The d⁶ Configuration: Iron (II) and Cobalt (III) are important species with this configuration. Iron (II) is quite stable although a mild reducing agent and cobalt (III) are stable in the presence of strong complexing reagents.

→ The d⁷ Configuration: The species with this configuration is cobalt (II) which is stable in aqueous solutions but gets oxidized to form CO (III) complexes in the presence of strong ligands.

→ The d⁸ Configuration: Nickel (II) is the most important species with this species.

→ The d⁹ Configuration: This configuration is found in Cu²⁺ compounds. It is by far the most important in the chemistry of copper.

→ The d¹⁰ Configuration: The two species Cu⁺ and Zn²⁺ are important with this configuration. Whereas Copper (I) is easily oxidized to copper (II), zinc (II) is the only state known for zinc.

General Group Trends in the Chemistry of the d-Block Metals:
Group 4: The titanium group of transition metals consists up of the elements titanium, zirconium and hafnium. The most important member of this group is titanium. It is extremely strong, has a high melting point and is resistant to corrosion. It is in great demand as a structural material. Zirconium and hafnium are silvery-white metals. The most stable oxidation state for the elements of this group is +4.

The titanium also possesses a + 3 oxidation state. The typical compounds of these elements are chlorides, TiCl₄, ZrCl₄, HfCl₄ and oxides TiO₂, ZrO₂ and HfO₂, ZrO₂ is a refractory material. Zirconium and hafnium occur together and exhibit similar properties.

Their atomic radii (Zr = 160 pm, Hf = 159 pm) are almost equal. This is due to the reason that usually increases in size down the group is cancelled by the lanthanide contraction. Zirconium and hafnium are both important for the generation of nuclear energy.

The vanadium group (Group 5): consists of vanadium, niobium and tantalum. The most stable oxidation state for this group is + 5. Vanadium is used as an additive to steel. The most important compound of vanadium is its pentoxide, V₂O₅, which is used as a catalyst in many reactions. Niobium alloys are used in jet engines, Tantalum is very resistant to corrosion and is used for making apparatus in chemical plants. It is also used in surgery, as for bone pins.

The chromium group (Group 6): contains the elements, chromium, molybdenum and tungsten. The most important oxidation states for chromium are + 3 and + 6, and for molybdenum and tungsten are + 5 and + 6. The zero oxidation states for these elements arise in metal carbonyls such as Cr (CO)₆.

These elements have very high melting and boiling points. Tungsten is the metal with the highest melting point. These metals are also very hard. Chromium is unreactive or passive at low temperatures because of the formation of a surface coating of oxide. It is due to this passive behaviour that chromium is used for electroplating iron to prevent rusting.

The metals of this group are very useful. Chromium is used in many ferrous alloys such as stainless steel, chrome steel, etc. It is also used for electroplating iron or some other metals to prevent corrosion. K₂Cr₂O₇ is a very important and useful compound of chromium.

Molybdenum and tungsten form very hard alloys with steel and are used in making cutting tools. Tungsten is also used as filament in electric bulbs. Molybdenum disulphide, MoS2 acts as a lubricant because it has a layer lattice.

Group 7: the manganese group: consist of the elements manganese, technetium and rhenium. These elements exhibit all the oxidation states from 0 to + 7, the most important being + 2, + 4 and + 7 for manganese and + 4 and + 7 for technetium and rhenium.

The elements of this group have quite a high melting and boiling points.

Manganese is obtained from its oxide ores by reduction with carbon or aluminium. Manganese metal does not have any use as such but is used in the manufacture of alloys such as ferromanganese (Fe + Mn) and manganese bronze (Mn + Cu + Zn). KMnO₄ is an important compound of Mn which is used as an oxidizing agent and finds so many other applications. Manganese dioxide is used as a catalyst. Rhenium is used in electronic filaments, high-temperature thermocouples and in flashbulbs.

The metals of group 8,9 and 10 are known as iron group metals: These elements are:

The first triad comprising of iron, cobalt and nickel is known as ferrous metals. These metals are ferromagnetic. Iron and cobalt exhibit oxidation states of + 3 and + 2 in their compounds while nickel compounds are generally in the + 2 oxidation state.

The elements of the second and third triad, ruthenium, rhodium, palladium, osmium, iridium and platinum are collectively known as platinum metals. These elements are relatively less abundant and exhibit a wider range of oxidation states. They are inert and serve as good catalysts.

The copper group (group 11): includes the elements copper, silver and gold. These metals are known as coinage metals. They form alloys with many metals. The most stable oxidation state for copper is + 2. For silver and gold, the oxidation state of + 1 is relatively more stable. The metals of this group have the highest electrical and thermal conductivities.

The zinc group (group 12): consists of the elements zinc, cadmium and mercury. The elements of this group show none of the characteristic properties of transition metals. The metals of this group are moderately electropositive and exhibit an oxidation state of + 2. Since their ionization energy is very high none of these metals possesses an oxidation state higher than + 2. Mercury, due to metal-metal bond also shows a formal oxidation state of +1. The elements of this group are diamagnetic. Mercury is the only metal that exists as a liquid.

Potassium Dichromate (K2Cr2O7):
Preparation: Chromite ore is fused with sodium or potassium carbonate in free access of air when the following reaction occurs.

The yellow solution of sodium chromate (Na₂CrO₄) is filtered and acidified with a calculated amount of sulphuric acid from which orange sodium dichromate Na₂Cr₂O₇. 2H₂O can be crystallised.
2 Na₂CrO₄ + 2H⁺ → Na₂Cr₂O₇ + 2Na⁺ + H₂O

Potassium dichromate is prepared by treating the solution of sodium dichromate with a calculated amount of potassium chloride.
Na₂Cr₂O₇ + 2 KCl → K₂Cr₂O₇ + 2 NaCl.

Orange crystals of potassium dichromate crystallise out. The dichromate and chromate are interconvertible in an aqueous solution depending upon the pH of the solution.
2 CrO₄^2- + 2 H⁺ → Cr₂O₇^2- + H₂O.
Cr2O₇^2- + 2 OH^– → 2 CrO₄^2- + H₂O.

The structures of chromate ionCrO₄^2-, and dichromate ion, Cr₂O₇^2- and shown below. Whereas CrO₄^2- is tetrahedral sharing one comer with Cr-O-Cr bond angle of 126°.

In acidic solution, both sodium and potassium dichromates are strong oxidizing agents, the former has a greater solubility in water.
Cr₂O₇^2- + 14 H⁺ + 6e^– → 2 Cr³⁺ + 7 H₂O [E° = 1.33 V]

→ Acidified solution oxidizes
1. iodides to iodine
6I^– → 3I₂ + 6e^–

2. sulphides to sulphur
3 H₂S → 6 H+ + 3 S + 6e^–

3. tin (II) to tin (IV)
3 Sn²⁺ → 3 Sn⁴⁺ + 6e^–

4. iron (II) salts to iron (III)
6 Fe²⁺ → 6 Fe³⁺ + 6e^–
The full ionic equation may be obtained by adding the two half-reactions. e.g.,
Cr₂07^2- + 14 H⁺ + 6 Fe²⁺ → 2 Cr³⁺ + 6 Fe³⁺ + 7 H₂O.

Potassium Permanganate, KMn04:
Preparation:
1. KMnO₄ is prepared by fusion of MnO, and KOH with an oxidizing agent like KNO₃.

Potassium manganate (K₂MnO₄) gives KMnO₄ in a neutral/acidic solution.
3 MnO₄^2- + 4H) → 2 MnO₄^– + MnO₂ + 2H₂O

2. Comfnercially, KMnO₄ is prepared from MnO₂ by electrolytic oxidation.

KMnO₄ forms dark purple crystals. It is not very soluble in water. When heated it decomposes at 240°C.

KMnO₄ shows weak temperature-dependent paramagnetism. It arises due to the coupling of the diamagnetic ground state of MnO₄ ion with paramagnetic excited states under the influence of the magnetic field.

The manganate and permanganate ions are tetrahedral. The green MnO₄^2- is paramagnetic with one unpaired electron, but MnO₄^– is diamagnetic.

Permanganic acid (HMnO₄) can be obtained by low-temperature evaporation of its aqueous solution. It is a strong oxidizing agent and in a pure state, it is explosive above 0°C.

→ Acidified KMnO₄ is a strong oxidizing agent:
MnO₄^– + e^– → MnO₄^2-
[Reduction half reaction]; E° = + 0.56 V

MnO₄^– + 4H⁺ + 3e^– → MnO₂ + 2H₂O
[Reduction half reaction]; E° = + 1.69 V.

MnO₄^– + 8H⁺ + 5e^– → Mn²⁺ + 4 H₂O
[Reduction half reaction]; E° = + 1.52 V.

A few important oxidizing reactions of KMnO₄ are given below.
1. In acid solutions:
(a) Iodine is liberated from potassium iodide:
10I^– + 2MnO₄^– + 16H⁺ → 2Mn²⁺ + 8H₂0 + 5I₂

(b) Fe²⁺ ion is converted to Fe³⁺:
5Fe²⁺ + MnO₄^– + 8H⁺ → Mn²⁺ + 4H₂O + 5Fe³⁺
The green iron (II) solution changes to yellow iron (III) solution.

(c) Oxalate ion or oxalic acid is oxidized at 60°C:
5 C2O₄^2- + 2MnO₄^–+ 16H⁺ → 2Mn²⁺ + 8H₂O + 10CO₂

(d) Hydrogen sulphide is oxidized, sulphur being precipitated:
H₂S + 2H⁺ → + S^2-
5S^2- + 2MnO₄^– + 16H⁺ → 2Mn²⁺ + 8H₂O + 5S

(e) Sulphurous acid or sulphite is oxidized to a sulphate or sulphuric acid:
5 SO₃^2- + 2MnO₄^– + 6H⁺ → 2Mn²⁺ + 3H₂O + 5SO₄^2-

(f) Nitrite is oxidized to nitrate:
5NO₂^– + 2MnO₄^– + 6H⁺ → 2Mn²⁺ + 5NO₃^– + 3H₂O

2. In neutral or faintly alkaline solution:
(a) A notable reaction is the oxidation of iodide to iodate:
2MnO₄^– + H₂O + I^– → 2MnO₂ + 2OH^– + IO₃^–

(b) Thiosulphate: s oxidised almost quantitatively to sulphate, a trace of thionate being formed
8MnO₄^– + 3S₂O₃^2- + H₂O → 8MnO₂ + 6SO₄^2- + 2OH^–

(c) Manganous salt is oxidised to MnOz; the presence of zinc sulphate or zinc oxide catalyses the oxidation:
2MnO₄^– + 3Mn²⁺ + 2H₂O → 5MnO₂ + 4H⁺

Note: Permanganate titrations in presence of hydrochloric acid are unsatisfactory since hydrochloric acid is oxidised to chlorine.

Uses: Besides its use in analytical chemistry, potassium permanganate is used as a favourite oxidant in preparative organic chemistry. Its uses for the bleaching of wool, cotton, silk and other textile fibres and for the decolourisation of oils are also dependent on its strong oxidising power.

→ The Inner Transition Elements (f-Block): It consists up of two series of elements.
1. Lanthanoids (58-71) which follow lanthanum (57) in the periodic table. are fourteen in number. Because La (57) closely resembles the lanthanoids, it is usually included in the study of lanthanoids and the general symbol Ln is often used.

2. Actinoids [the fourteen elements (90-103)] which follow actinium (89). A discussion of actinoids also includes actinium (Ac), besides the 14 elements.

In the case of lanthanoids, 4f orbitals/are filled up and in the case of actinoids, 5f orbitals are filled up. That is why they are called Inner Transition Elements as the last electron (also called differentiating electron) enters the antepenultimate energy level, ie.., (n – 2) f-orbitals: inner to the penultimate energy level and they form a transition series within the transition series (d-block elements)

Lanthanides or Landhanoids or Lanthanoness: In these elements, the last electron enters the 4f-orbitals and is also referred to as the first inner transition series. Earlier they were called rare earth.

Actinides or Actinoids or Actions: In these elements, the last electron enters one of the 5f-orbitals and is also required as the second inner transition series.

The study of lanthanoids is easier because they show only one stable oxidation state (+ 3). On the other hand, the chemistry of actinoids is much more complicated partly because they show a wide range of oxidation states and partly because they are radioactive.

→ Electronic Configuration: General E.C. of lanthanoids and actinoids is (n – 2) f^1-14 (n – 1) d^0-1 ns². Thus they have three incomplete shells, viz., (n – 2), (n – 1) and nth. Their electronic config. is given below.

Trends in Ionic Radii of Trivalent Lanthanoids

Table: Electronic configurations and radii of lanthanum and lanthanoids:

Only electrons outside [Xe] core are indicated.

The overall decrease in atomic and ionic radii from lanthanum to lutetium (the lanthanoid contraction) is a unique feature in the tire chemistry of lanthanoids. It is due to the imperfect shielding of one 4f electron from another 4f electron due to the highly diffused shape of f-orbitals. With the increase in nuclear charge, there is a fairly regular decrease in their sizes.

The cumulative effect of the contraction of the lanthanoid series called lanthanoid contraction causes the radii of the third transition series to be very similar to those of the correspond ing members of the 2nd series. As a.result Zr (160 pm) and Hf (159 pm) have identical radii and so exist together.

Thus Zr and Hf face difficulty in their separation.
Colour and Paramagnetism: Many trivalent lanthanoids are coloured both in the solid-state and in aqueous solutions. It is due to the presence of f-electrons. Neither Lu³⁺ ion (f°) nor Lu³⁺ (f¹⁴) shows any colour. All others (having f¹ to f¹³ arrangement) show colour. The lanthanoid ions other than f° (La³⁺, Ce⁴⁺) type and the f¹⁴ type (Yb²⁺, Lu³⁺) are all paramagnetic due to the presence of unpaired electrons. Paramagnetism is maximum in neodymium.

→ Ionisation Enthalpies: The IE₁ and IE₂ are around 600 kJ mol^-1 and 1200 kJ moH comparable with those of calcium. The values of IE₃ indicate that the exchange energy considerations impart a degree of stability to empty, half-filled and completely filled f-level (f°, f⁷, f¹⁴). It is evident from the abnormally low values of IE₃ of lanthanum, gadolinium and lutetium.

→ Oxidation States: All the lanthanoids show a + 3 oxidation state which is most significant: Some of them show + 2 and + 4 oxidation states also as ions in solution or in solid compounds. This irregularity arises mainly from the extra stability of empty, half-filled or completely filled f-subshell. Ce^IV [₁₆Rn-noble gas] is formed easily, but it is a strong oxidant reverting to the common oxidation state of + 3.
Ce⁴⁺ + e^– → Ce³⁺] Reduction half reaction
E° for the above is = + 1.74 V

If can oxidise water, but the reaction rate is slow. Pr, Nd, Tb and Dy also exhibit + 4 states in oxides MOr Eu²⁺ is formed by losing the two s-electrons (F). However, Eu²⁺ is a strong reducing agent changing to the common oxidation state + 3.
EU²⁺ → EU³⁺ + e^–

Similarly Yb²⁺ [f¹⁴] is a reductant. Tb (IV) [f⁷] is an oxidant.
Tb4+ + e- → Tb3+

Samarium also shows oxidation states of + 2 and + 3 like Eu.

Properties and Uses:
1. All the lanthanides are silvery-white soft metals and tarnish rapidly in the air. Hardness increases with increasing atomic number, Samarium is steel hard.

2. M.Pts are in the range of 1000 to 1200 K, but Sm melts at 1623 K

3. They have a metallic structure.

4. They are good conductor of heat and electricity.

5. In their chemical behaviour, lanthanoids are generally similar to calcium (Ca = 20), but with increasing atomic number, they behave more like aluminium.

6. Values for E° for the reduction half-reaction.
Ln³⁺ (aq) + 3e^– → Ln (s) are in the range of – 2.2 to – 2.4 V except for Ln = Eu, for which it is – 2.0 V,

7. The metals combine with hydrogen when heated gently.
2 Ln + 3 H₂ → 2 Ln H₃

8. When heated in carbon, carbides of the type Ln₃C, Ln₂C₃ and LnC₂ are formed.

9. They liberate H2 gas from dilute acids.
2 Ln + 6HCl (dil.) → 2 LnCl₃ + 3H₂I

10. They bum in halogens to form halides.
2 Ln + 3X₂ → 2 LnX₃

11. They form oxides and hydroxides of the type M₂O₃ and M (OH)₃. They are basically like alkaline earth metal oxides and hydroxides. A summary of the chemical reactions of lanthanoids are given below:

The best use of the lanthanoids is for the production of alloy steels for plates and pipes. A well-known alloy is Misch Metall which consists of a lanthanoid metal (~ 95%) and iron (~ 5%) and traces of S, C, Ca and Al. It is used to produce bullets, shell and lighter flint. Mixed oxides are used as catalysts in cracking. Some Ln oxides are used as phosphors in TV screens.

→ The Actinoids: The elements which follow actinium (89) in the periodic table are called actinoids [Starting from Th – 90 to Lr – 103.]

→ Actinoids are radioactive elements. Earlier members have relatively long half-lives and later ones from day to 3 minutes for Lr. This makes their study difficult.

→ Electronic Configuration: All the actinoids have a 7s² electron configuration and variable occupancy of the 5f and 6d subshells. The irregularities in the electronic configuration of actinoids like those in lanthanoids are related to the stabilities of 5f°, 5f⁷ and 5f¹⁴. 5f orbitals participate in bonding to a greater extent.

→ Oxidation states: There is a greater range of oxidation states shown by actinoids due to the comparable energies of 5f, 6d and 7s levels.

The electronic configurations and oxidation states of actinoids are in the following tables.

Table: Oxidation states of actinium and actinoids:

Unlike 4f orbitals of lanthanoids which are buried, 5f orbitals of actinoids participate in bonding to a far greater extent. In addition to showing an oxidation state of + 3 like lanthanoids, actinoids also show oxidation states of + 4, + 5, + 6 and + 7. Actinoids which show several oxidation states [Np, Pu, Am] make it difficult to review their chemistry.

Physical and Chemical reactivity of actinoids

• Actinoids are all silvery-white metals.
• They are highly reactive metals, especially when finely divided.
• Boiling water reacts with them to give a mixture of oxides, hydrides.

Some applications of d-Block elements

1. Iron and Steel are the most important construction materials.
2. Compounds like TiO are prepared for use in the pigment industry.
3. Mn02 is used in dry battery cells. The battery industry also requires Zn and Ni/Cd.
4. Cu, Ag and Au are coinage metals.
5. Many of the metals and their compounds find use as catalysts like finely divided Ni, V₂O₅, TiCl₄, Fe etc.
6. PdCl₂ is used in the Wacker process for the oxidation of ethyne to ethanol
7. Ni-complexes are used in the polymerisation of alkynes and other organic compounds like benzene.
8. AgBr is used in the photographic industry.

By going through these CBSE Class 12 Chemistry Notes Chapter 7 The p-Block Elements, students can recall all the concepts quickly.

The p-Block Elements Notes Class 12 Chemistry Chapter 7

The p-block elements are placed in groups 13 to 18 of the periodic table. Their valence shell electronic configuration is ns²np⁶ to ns²np⁶ (except He which has 1s² electronic configuration.)

Group 15 Elements: Group 15 elements include Nitrogen (N), Phosphorus (P), Arsenic (As), Antimony (Sb) and Bismuth (Bi).

Table 7.1: Atomic and Physical Properties of Group 15 Elements

^αE₁₁₁ single bond (E = element); ^bE^3-; ^cE³⁺; ^dWhite phosphorus; ^eGrey α-format 38.6 atm of Sublimation temperature; g At 63 K h Grey α-form; Molecular N₂

As we go down the group, there is a shift from non-metallic to the metallic character through metalloids character. N and P are non-metals, AS and Sb are metalloids and Bi is a typical metal.

→ Occurrence: N₂ comprises 78% by volume of the atmosphere. In the earth’s crust, it occurs as sodium nitrate (NaNO₃) called Chile saltpetre and potassium nitrate (KNO₃) called Indian saltpetre. It is found in the form of proteins in plants and animals. Phosphorus (P) occurs in minerals of the apatite family Ca₉ (PO₄)₆. CaX₂ [X = F, Cl, or OH] e.g. fluorapatite Ca₉ (PO₄)₆. CaF₂ which are the main components of the phosphate rocks. P is an important constituent of animal and plant matter and is present in bones and living cells.

→ Electronic Configuration: The valence shell electronic configuration of these elements is ns2np3. The s orbital is completely filled and p orbitals are exactly half-filled (p_x¹ p_y¹ P_z¹). Thus their electronic configuration is extra stable.

→ Atomic and Ionic Radii: Covalent and ionic radii increase in size down the group due to the addition of the orbit each time. There is a considerable increase from N to P. However, from As to Bi only a small increase in covalent radius is observed. This is due to the presence of completely filled d and /or f orbitals in heavier members.

→ Ionization Enthalpy: I.E. values of Group 15 elements is much greater than that of Group 14 elements due to the extra-stability of half-filled p-orbitals and smaller size. It decreases from top to bottom within the group due to a gradual increase in atomic size. The order of successive ionisation enthalpies is Δ_iH₁ < Δ_iH₂ < Δ_iH₃.

→ Electronegativity: Electronegativity decrease down the group. The difference is not much pronounced amongst the heavier elements.

→ Physical Properties: All the elements are polyatomic. Dinitrogen (N₂) is a diatomic gas while others like P₄ etc. are solids. The b. pts; in general, increase from top to bottom in the group, but them. pts increase up to As and then decrease up to Bi.

Except for Nitrogen, all the elements show allotropy.

Chemical Properties:
Oxidation states and trends in chemical reactivity: Tire common oxidation states of these elements are – 3, + 3 and + 5. The tendency to exhibit – 3 oxidation state decreases down the group due to an increase in size and metallic character.

N does not show + 5 oxidation state due to the absence of d-orbitals and Bi does not show + 5 oxidation state due to the Inert pair Effect. The only well characterised Bi (V) compound is BiF₅. N exhibits + 1, + 2, + 4 oxidation states also when it reacts with oxygen. P also exhibits + 1 and + 4 oxidation states in some oxoacids.

In the case of N, all oxidation states from + 1 to + 4 tend to disproportionate into acid solution1. For example
3 HNO₂ → HNO₃ + 2 NO + H₂O

Similarly, in the case of P nearly all intermediate oxidation states disproportionate into + 5 and – 3 both in alkali and acid,

Anomalous properties of Nitrogen: N differs from its congeners due to

1. its smaller size.
2. high electronegativity and high ionisation enthalpy.
3. non-availability of d-orbitals.

Nitrogen has a unique ability to form an-pit multiple bonds with itself and with other elements having small size and having high electronegativity (e.g., C, O)

N₂ has a triple bond: N ≡ N [one a and two K bonds]
∴ Its bond enthalpy (941.4 kJ mol^-1) is very high.
P and As can form dπ-pπ bonds with transition metals.

1. Reactivity towards hydrogen: All the elements of Group 15 form hydrides of the type EH₃ where E = N, P, As, Sb, Bi. The stability of hydrides decreases from NH₃ to BiH₃ and hence reducing character increases from NH₃ to BiH₃. Basicity of hydrides also decreases in the order NH₃ > PH₃ > AsH₃ > SbH₃ > BiH₃.

Table 7.2: Properties of Hydrides of Group 15 Elements

2. Reactivity towards oxygen: All these elements form oxides: E₂O₃ and E₂O₅. The oxide in the higher oxidation state is more acidic.
Their acidic character decreases down in a group. The oxides of the type EO₃ of N and Pare purely acidic, that of As, Sb amphoteric and those of Bi predominantly basic.

3. Reactivity towards halogens:
2E + 3X₂ → EX₃
2E + 5X₂ → 2EX₅
N does not form pentahalide due to absence of d-orbitais. All trihalides are stable except that of N.

Pentahalides are more covalent than trihalides.

4. Reactivity towards metals: These elements combine with metals showing – 3 oxidation state in compounds like Ca₃N₂, Ca₃P₂, Mg₃Bi₂.

Dinitrogen, N₂:
Preparation:
1. Commercially dinitrogen is produced by the liquefaction and fractional distillation of air when N₂ distils the first.

2. Lab. Method:
NH₄Cl (aq) + NaNO₂ (aq) → N₂ (g) + 2H₂O (l) + NaCl (aq)

3. Pure N₂ can be obtained from sodium or barium azide.

→ Properties: Dinitrogen is a colourless, odourless, tasteless and non-toxic gas. It has two stable isotopes: ¹⁴N, ¹⁵N. It has very low solubility in water.

Because of the high bond enthalpy of N ≡ N, dinitrogen is rather inert at room temperature.

At high temperature, it directly combines with some metals to form largely ionic nitrides and with non-metals, covalent nitrides.

Haber’s process:

Uses:

1. Manufacture of NH₃ and Calcium Cyanamide.
2. To create an inert atmosphere.
3. Liquid N₂ is used as a refrigerant,

Ammonia:
Preparation:
1. From Urea:
NH₂CONH₂ + 2H₂O → (NH₄)₂ CO₃ → 2NH₃ + H₂O + CO₂

2. From ammonium salts and caustic soda/lime.
2NH₄Cl + Ca(OH)₂ → 2NH₃ + 2H₂O + CaCl₂
(NH₄)₂ SO₄ + 2 NaOH → 2NH₃ + 2H₂O + Na₂SO₄

3. On a large scale, it is prepared by Haber’s process.

The pressure is 200 × 10⁵ Pa and iron oxide with small amounts of K₂O and Al₂O₃ is used as a catalyst,

Properties of Ammonia:

1. It is a colourless gas with a pungent smell.
2. It shows hydrogen bonding in solid and liquid states.

Structure of NH3:

3. It is highly soluble in water. Its aqueous solution is weakly basic.
NH₃ (g) + H₂O (l) ⇌ NH₄⁺ (aq) + OH^– (aq)

4. It forms ammonium salts with acids.
NH₃ + HCl → NH₄Cl

5. As a weak base it precipitates the hydroxides of many metals
from their salt solutions, e.g.,

6. Due to its tendency to donate a pair of electrons on N, it acts as a Lewis-Base, forms complexes with Cu²⁺, Ag⁺.


Uses:

1. Production of nitrogenous fertilizers.
2. Production of nitric acid.
3. Liquid ammonia is a non-aqueous solvent.
4. Liquid NH₃ is a referigerant.

Oxides of Nitrogen: Nitrogen forms a number of oxides in different oxidation state. The details of names, formulae, ox. state, preparation, colour etc. are given:

Lewis dot main resonance structures and bond parameters of oxides are given on the next page.

Table 7.4 Structures of Oxides of Nitrogen

Nitric Acid, HNO₃
Preparation: It is prepared in the laboratory by heating NaNO₃ or KNO₃ and cone. H₂SO₄.

Method of manufacture of Nitric acid by Ostwald’s process.

Step:
1. Catalytical oxidation of NH₃ by O, or air.

2. NO thus formed combines with 02 giving NO₂
2NO (g) + O₂ (g) ⇌ 2NO₂(g).

3. N02 SO formed dissolves in water to give HNO₃.
3NO₂ + H₂O(l) → 2 HNO₃ + NO (g)

Structure of HNO₃: In the gaseous state, HNO₃ exists as a planar molecule.

In an aqueous solution, it behaves as a strong acid.
HNO₃ (aq) + H₂O (Z) → (aq) + NO₃ (aq).
Concentrated HNO₃ is a strong oxidizing agent.

→ Reaction with Cu:
3Cu + 8 HNO₃ (dil) → 3 Cu (NO₃)₂ + 2 NO + 4 H₂O
Cu + 4 HNO₃ (cone.) → Cu (NO₃)₂ + 2NO₂ + 2H₂O

→ Reaction With Zinc
4 Zn + 10 HNO₃ (dil.) → 4 Zn (NO₃)₂ + 5 H₂O + N₂O
Zn + 4HNO₃ (cone.) → Zn (NO₃)₂ + 2H₂O + 2NO₂

→ Reaction with I₂

→ Reaction with Carbon
C + 4 HNO₃ → CO₂ + 2H₂O + 4NO₂

→ Reaction with Sulphur
S₈ + 48 HNO₃ (Cone.) → 8H₂SO₄ + 48 NO₂ + 12 H₂O

→ Reaction with Phosphorus
P₄ + 20 HNO₃ → 4H₃PO₄ + 20 NO₂ + 4H₂O

→ Brown Ring Test for nitrates (NO₃^–)
NO₃– + 3 Fe²⁺ + 4 H⁺ → NO + 3 Fe³⁺ + 2H₂O.

Uses of Nitric acid:

1. Oxidant in the laboratory.
2. manufacture of ammonium nitrate for fertilizers.
3. manufacture of explosive like TNT, nitroglycerine. Phosphorus: Phosphorus exists in three important allotropic forms: white, red and black.

White Phosphorus: White Phosphorus consists of discrete tetrahedral molecules as shown:

White phosphorus P₄

Red Phosphorus: Red Phosphorus is obtained by heating white phosphorus at 573 K in an inert atmosphere for several days. It is much less reactive than white phosphorus. It is polymeric, consisting of P4 tetrahedral linked together as shown in Fig. below.

Red Phosphorus

Black Phosphorus has two forms α-black phosphorus and β-black phosphorus. α-Black phosphorus is formed when red phosphorus is heated in a sealed tube at 803 K. It can be sublimed in air and has opaque monoclinic or rhombohedral crystals. It does not oxidise in the air. β-Black phosphorus is prepared by heating white phosphorus at 473 K under high pressure. It does not bum in the air up to 673 K.

Phosphine (PH₃) Preparation:
1. From calcium phosphide and H₂O/dil. HCl
Ca₃P₂ + 6H₂O → 3Ca(OH)₂ + 2PH₃.
Ca₃P₂ + 6HCl → 3CaCl₂ + 2PH₃.

2. Lab. Method: By heating white phosphorus with cone. NaOH in an inert atom of CO₂.
P₄ + 3NaOH +- 3H₂O → PH₃ + NaH₂PO₂ (Sod. hypophosphite)

Properties:

1. Colourless gas with rotten fish smell, highly poisonous.
2. Explodes in contact with traces of oxidizing agents like HNO₃ Cl₂ etc.,
3. Slightly soluble in water.
4. When absorbed in CuSO₄ or HgCl₂ solutions, phosphides are obtained.
   3 CuSO4 + 2 PH₃ → Cu₃P₂ + 3H₂SO₄
   3HgCl₂ + 2PH₂  → Hg₂ P₂ + 6HCl.
5. PH₃ is weakly basic like NH₃.
   PH₃ + HBr → PH₄Br.

Phosphorus halides: P forms two types of halides, PX₃ [X = F, Cl, Br, I]and PX₅[X = F,Cl,Br]
Phosphorus trichloride (PCl₃)

Preparation:
P₄ + 6Cl₂ → 4PCl₃.

Properties:

1. Colourless oily liquid.
2. Hydrolysis in the presence of mixture.
   PCl₃ + 3H₂O → H₃PO₃ + 3HCl

Shape: It has a pyramidal shape as shown in which phosphorus is sp³ hybridized.

Shape of PCl₃

Phosphorus pentachloride (PCl5)
Preparation:

Properties:

1. It is a yellowish-white powder.
2. In moist air, it hydrolyses first to PoCl₃, and then H₃PO₄.
   PCl₅ + H₂O → POCl₃ + 2 HCl
   POCl₃ + 3H₂O → 4 H₃PO₄ + 3HCl
3. On heating, it sublimes but decomposes on stronger heating.
   

In gaseous and liquid phases, it has a trigonal bipyramidal structure shown below. The three equatorial P-Cl bonds are equivalent, while the two axial bonds are longer than equatorial bonds. This is due to the fact that the axial bond pairs suffer more repulsion as compared to equatorial bond pairs.

In the solid-state, it exits as an ionic solid, [PCl₄]⁺[PCl₆]^– in which the cation, [PCl₄]⁺ is tetrahedral and the anion, [PCl₆]⁺ octahedral.

Oxo-Acids of Phosphorus: P forms a number of oxo-acids like

1. Hypophosphorous acid (H₃PO₂) (ox. state + 1)
2. Orthophosphoric acid (H₃PO₃) (ox. state + 3)
3. Pyrophosphrus acid (H₄P₂O₅) (ox. state + 3)
4. Hypophosphoric acid (H₄P₂O₆) (ox. state +4)
5. Orthophosph’oric acid (H₃PO₄) (ox. state + 5)
6. Pyrophosphoric acid (H₄P₂O₇) (ox. state +5)
7. Metaphosphoric acid (HPO₃) (ox. state + 5)

The compositions of the oxo-acids are interrelated in terms of loss or gain of H₂O molecule or oxygen atom.
These acids in the + 3 oxidation state of P tend to disproportionate as follows:

The acids which contain the P-H bond are strong reducing agents.

Group 16 Elements: Oxygen (O), Sulphur (S), Selenium (Se), Tellurium (Te) and Polonium (Po) constitute Group 16 of the periodic table. They are also called Chalcogens (ore-forming).

→ Occurrence: Oxygen is the most abundant of all the elements of the earth. Dry air contains 20.95% oxygen by volume. However Sulphur is present in the earth’s crust to 0.03 – 0.1% only. Combined Sulphur is present in gypsum CaSO₄. 2H₂O, Epsom salt MgSO₄.7H₂O. Selenium and tellurium are also found as metal selenides and tellurides in sulphide ores.

→ Electronic Configuration: The elements of Group 16 have six electrons in the outermost shell and have ns2np4 general electronic configuration.

→ Atomic and Ionic Radii: Due to the increase in the no. of shells, atomic and ionic radii increase from top to bottom, The size of the oxygen atom is, however, exceptionally small.

→ Ionisation Enthalpy: Ionisation enthalpy decreases down the group.

Table: Some Physical Properties of Group 16 Elements

^aSingle bond; ^bApproximate value; ^cAt the melting point; ^dRhombic sulphur; ^eHexagonal grey; ^fMonoclinic form, 673 K.
Oxygen shows oxidation states of + 2 and +1 in oxygen fluorides OF₂ and O₂F₂ respectively.

→ Electron Gain Enthalpy: Because of the compact nature of the oxygen atom, it has less negative electron gain enthalpy than sulphur.

→ Electronegativity: Next to F, oxygen has the maximum value of electronegativity value amongst the elements. Within the group, its value decreases from top to bottom implying that metallic character increases from O to Po.

→ Physical Properties: Oxygen and Sulphur are non-metals, Selenium and Tellurium metalloids, whereas Polonium is a metal. Pollonium is radioactive [t_1/2 = 13.8 days]

All these elements exhibit allotropy.
M.Pts and B.Pts. increase with the increase in atomic no. down the group. The large difference between the M.Pts and B.Pts. of oxygen and sulphur may be explained on the basis of their atomicity: Oxygen exists as diatomic molecules (O₂) whereas Sulphur exists as polyatomic (S₈).

Chemical Properties:
Oxidation states and trends in chemical reactivity:
O = – 2, -1, +1, (in O₂F₂), + 2 (in OF₂)
S = – 2, + 2, +4, + 6 (in SF₆)
Se = – 2, (+ 2), +4, + 6 (SeF₆)
Te=-2,(+2), +4, + 6(TeF₆)
Po = + 2, + 4
[Within brackets less stable.]

The stability of + 4 oxidation state increases and that of + 6 decreases due to the Inert-pair Effect. Bonding in + 4, + 6 oxidation states are primarily covalent.

The anomalous behaviour of oxygen is due to its

1. small size,
2. high electronegativity,
3. absence of d-orbitals.

1. Reactivity with Hydrogen.
(a) There is strong hydrogen bonding in H₂O which is not found in H₂S.
(b) All the elements of Group 16 form hydrides of the type H₂E[E = S, Se, Te, Po]
(c) H₂O is neutral. Other hydrides are acidic and acidic character increases.
H₂O < H₂S < H₂Se < H₂Te < H₂Po
(d) Thermal stability decreases:
H₂O > H₂S > H₂Se > H₂Te > H₂Po
(e) H₂O does do not have reducing character. Reducing character increases from H₂S < H₂Se < H₂Te.

2. Reactivity with Oxygen: All these elements form oxides of EO₂ and EO₃ types [E = S, Se, Te, or Po]. Both types are acidic in nature.

3. Reactivity towards halogens: Elements of Group 16 form a large no. of halides of the type, EX₆, EX₄, EX₂ where X = halogen. The stability of the halides decreases in the order F^– > Cl > Br > I^–

Dioxygen, O₂:

Industrially, O₂ is obtained from air by first removing CO₂ and water vapour and then, the remaining gases are liquified and fractionally distilled to give N₂ and O₂.

Properties of O₂:

1. Colourless, odourless gas.
2. 30.8 cm³3 of O₂ soluble per litre of water at 293 K.
3. It has three isotopes ¹⁶O, ¹⁷O, ¹⁸O.
4. Molecular oxygen (O₂) is paramagnetic.
5. Reactions with metals, non-metals and other compounds are as follows:
   
   

Uses:

• It is used in oxyacetylene welding.
• Manufacture of many metals, particularly steel.
• Oxygen cylinders are widely used in hospitals, high altitude flying and mountaineering.
• The contraction of fuels, e.g., hydrazines in liquid oxygen, provides tremendous thrust in rockets.

→ Simple Oxides: A binary compound of oxygen with another element is called oxide.

Oxides can be simple like MgO, Al₂O₃ or mixed like Pb₃O₄, Pe₃O₄.
In general metallic oxides like CaO, BaO, Na₂O are basic. Non-metal oxides like SO₂, CO₂ are acidic.

Some oxides like Al₂O₃ are amphoteric. They react with acids as well as bases.
Al₂O₃ + 6 HCl → 2AlCl₃ + 3H₂O.
Al₂O₃ + 2NaOH → 2Na AlO₂ + H₂O

Ozone, O₃: Ozone is an allotropic form of oxygen. At a height of 20 km from sear level, it is formed from atmospheric oxygen in the presence of sunlight.

The ozone layer protects the earth’s surface from UV rays.

Preparation:

ΔH° = + 142 kJ mol^-1

Properties:

1. It is a pale blue gas, dark blue liquid and violet-black solid.
2. O₃ is thermodynamically unstable w.r.t. O₂.
   ∴ It acts as a powerful oxidizing agent.

O₃ → O₂ + O
1.  It oxidizes lead sulphide to lead sulphate.
PbS + 4O₃(g) → PbSO₄ (s) + 4O₂

2. It oxidizes I^– ions to iodine
2I^– (aq) + H₂O (l) + O₃ (g) → 2 OH^– (aq) + I₂ (s) + O₂ (g)

3. It oxidizes NO to NO₂
NO (g) + O₃ (g) → NO₂ (g) + O₂ (g)

Structure: The two oxygen-oxygen bond lengths in the ozone molecule are identical (128 pm) and the molecule is angular as expected with a bond angle of about 117°. It is a resonance hybrid of two main forms:

Uses:

1. It is used as a germicide, disinfectant and for sterilising water.
2. It is also used for bleaching oils, ivory, flour, starch.
3. It is an oxidizing agent.

Sulphur-Allotropic Forms: Out of numerous allotropes, the two most important forms are:

• Yellow rhombic (α-sulphur),
• Monoclinic (β-Sulphur)
  

The temperature 369 K is called Transition temperature.

The structure of (a) S₈ ring in rhombic sulphur and (b) S₆ form:

Both rhombic and monoClinic Sulphur have S₈ molecules. These S₈ molecules are packed to give different crystal structures. The S₈ ring in both forms is puckered and has a crown shape.

In cyclo-S₆ the ring adopts the chair form. At elevated temperatures (~ 1000 K) S₂ is the dominant species and is paramagnetic like O₂

Sulphur Dioxide (SO₂):
Preparation:
1. By burning S in air or O₂.
S(s) + O₂(g) → SO₂(g)

2. Lab. Method: By treating a sulphite with dilute H₂SO₄4.
SO₃^2- (aq) + 2H⁺ (aq) → H₂O (l) + SO₂(g)

3. Industrially, it is produced by roasting of sulphide ores:

Properties:
1. It is a colourless gas with a pungent smell.

2. It is highly soluble in water.

3. It liquefies at room temp, under a pressure of 2 atmospheres.

4. Liquified SO₂ boils at 263 K.

5, SO₂ when dissolved in water, forms sulphurous acid.
SO₂(g) + H₂O (l) → H₂SO₃ (aq)

6. Reaction with NaOH:


7. Reaction with Cl₂:

8. Reaction with O₂:

9. Moist SO₂ behaves as a reducing agent. It reduces iron (III) ions to iron (II) ions and decolourised acidified potassium permanganate(VII) solution. It is a test for SO₂ gas.
2Fe³⁺ (aq) +SO₂ + 2H₂O → 2 Fe²⁺ (aq) + SO₄^2- + 4H⁺
5 SO₂ + 2 MnO₄^– + 2H₂O → 5 SO₄ ^2- + 4H⁺ + 2 Mn²⁺

Structure: The molecule of SO₂ is angular. It is a resonance hybrid of the two canonical structures

Uses:

• Refining of petroleum, sugar
• Bleaching wool and silk.
• Antichlor, disinfectant and preservative.
• Manufacture of H2S04 and other industrial chemicals.
• Liquid S02 is used as a solvent.

Oxo-Acids of Sulphur: Sulphur forms a number of oxo-acids such as H₂SO₂, H₂S₂O₃, H₂S₂O₄, H₂S₂O₅, H₂S_xO₆ (X = 2 to 5), H₂SO₄, H₂S₂O₇, H₂SO₅, H₂S₂O₈.

Structures of Sulpiwrous and Sulphuric acids

Sulphuric Acid (H₂SO₄): It is called the king of chemicals and is one of the most widely used industrial chemicals.

→ Manufacture: Sulphuric acid is manufactured by Contact Process which involves 3 steps:

1. Burning of Sulphur or Sulphide ores in the air to produce SO₂.
2. Conversion of SO₂ into SO₃ by reaction with air in the presence of catalyst V₂O₅.
3. Absorption of SO₃, in H₂SO_4, to give Oleum (H₂S₂O₇).
   The key step is the catalytical oxidation of SO₂ with O₂ to give SO₃ in the presence of V₂O₅.
   
   The reaction is exothermic, reversible and proceeds with a decrease in volume.
   ∴ Low temperature (720 K) and high pressure (2 bar) are favourable conditions for maximum yield.

The tempi should not be very low otherwise the rate of reaction will below.

The SO₃ gas from the catalytical converter is absorbed in concentrated H₂SO₄ to produced Oleum (H₂S₂O₇) which on dilution with water gives H₂SO₄ of the desired concentration.
SO₃ + H₂SO₄ → H₂S₂O₇ (oleum)
H₂S₂O₇ + H₂O → 2 H₂SO₄

The sulphuric acid obtained by the contact process is 96-98% pure.

Flow diagram the iiai1ufat true of sulphuric acid

Properties:

1. H₂SO₄ is a colourless, dense, oily liquid.
2. It dissolves in water producing a large amount of heat. Hence it should be slowly added to water (and never water in acid)

The chemical properties of H₂SO₄ are due to its
(a) low volatility,
(b) strong acidic character,
(c) strong affinity for water,
(d) ability to act as an oxidizing agent.

In water, it ionises in 2 steps:
1. H₂SO₄ (aq) + H₂O (l) → H₃O+ (aq) + HSO₄^– (aq)
K_a1 = very large (> 10)

2. HSO₄^–(aq) + H₂O (l) → H₃O⁺ (aq) + SO₄ ^2- (aq);
K_a2 = 1.2 × 10^-2.

(a) It is used to prepare more volatile acids:
NaNO₃ + H₂SO4 → NaHSO₄ + HNO₃.
2 NaCl + H₂SO₄ → Na₂SO₄ + 2 HCl

(b) Cone. H₂SO₄ is a strong dehydrating agent

(c) Hot cone. H₂SO₄ is a strong oxidising agent
H₂SO₄ → H₂O + SO₂ + O

(d) Reaction with metals and non-metals:
Cu + 2 H₂SO₄ → CuSO₄ + SO₂ + 2H₂O
C + 2H₂SO₄ (cone.) → CO₂ + 2SO₂ + 2H₂O

Uses of H₂SO₂ :

• It is used in the manufacture of hundreds of other compounds and also in many industrial processes.
• Manufacture of fertilizers.
• Petroleum refining and manufacture of pigments, paints and dyestuff intermediates.
• Detergent industry.
• Metallurgical applications.
• Storage batteries and so on.

Group 17 Elements: Fluorine, chlorine, .bromine, iodine and astatine are the members of Group 17. They are collectively called halogens (sea-salt producers)

→ Occurrence: Fluorine is present as CaF₂ (fluorspar) and Na₃ AlF₆ (cryolite). Seawater contains chlorides, bromides and iodides of sodium, potassium, magnesium and calcium. The deposits of dried up seas contain NaCl and carnallite, KCl. MgCl₂,6Fl₂O. Seaweeds contain 0.5% of iodine and Chile saltpetre contains up to 0.2% of sodium iodate.

→ Electronic Configuration: All these elements have seven electrons in their valence shell [ns²np³] which is 1 electron less than the next noble gas.

→ Atomic and ionic radii: The halogens have the smallest atomic radii in their respective periods due to the maximum effective nuclear charge. The size of the anions is greater than the corresponding atoms. The atomic and ionic radii increase from top to bottom.

→ Ionisation Enthalpy: They have little tendency to lose electrons. Thus they have very high ionisation enthalpy. I.E. decreases from top to bottom.

→ Electronegativity: They have high electronegativities. It decreases from top to bottom. Fluorine is the most electronegative element in the periodic table. [F = 4.0]

→ Electron Gain Enthalpy: Halogens have maximum negative- electron gain enthalpy in their corresponding periods. It becomes less negative down the group. Electron gain enthalpy of F is less than that of Cl due to the compact size of F.

→ Physical Properties: F₂ and Cl₂ are gases. Br₂ is a liquid and I₂ is solid. They are all diatomic. M.pts and B.pts increase from top to bottom. F₂ is yellow, Cl₂ is greenish-yellow, Br₂ is red and I₂ violet in colour.

F₂ has a smaller enthalpy of dissociation as compared to Cl₂.

→ Chemical Properties:
Oxidation States and Trends in Chemical Reactivity: All the halogens exhibit a -1 oxidation state. However, chlorine, bromine and iodine exhibit +1, + 3, +5 and +7 oxidation states also as explained below:

Table: Atomic and Physical Properties of Halogens (Group 17)

^aRadioactive; ^bPauling scale; ^cFor the liquid at temperatures (K) given in the parentheses; ^dsolid; ^eThe half-cell reaction is X₂(g) + 2e^– → 2X^–(aq).

The higher oxidation states of chlorine, bromine and iodine are realised mainly when the halogens are in combination with the small and highly electronegative fluorine and oxygen atoms, e.g., in interhalogens, oxides and oxoacids.

The oxidation states of +4 and +6 occur in the oxides and oxoacids of chlorine and bromine. The fluorine atom has no d-orbitals in its valence shell and therefore cannot expand its octet. Being the most electronegative, it exhibits only a -1 oxidation state.

All the halogens are highly reactive. They react with metals and non-metals to form halides. The reactivity of the halogens decreases down the group.

The ready acceptance of an electron is the reason for the strong oxidising nature of halogens. F2 is the strongest oxidising halogen and it oxidises other halide ions in solution or even in the solid phase.

In general, a halogen oxidises halide ions of higher atomic number.
F₂ + 2X^– → 2F + X₂ (X = Cl, Br or I)
Cl^– + 2X^– → 2CT + X₂ (X = Br or I)
Br₂ + 2I^– → 2Br + I₂

The decreasing oxidising ability of the halogens in aqueous solution down the group is evident from their standard electrode potentials which are dependent on the parameters indicated below:

The relative oxidising power of halogens can further be illustrated by their reactions with water. Fluorine oxidises water to oxygen whereas chlorine and bromine react with water to form corresponding hydrohalic and hypohalous acids.

The reaction of iodine with water is non-spontaneous. In fact, I^– can be oxidised by oxygen in an acidic medium; just the reverse of the reaction observed with fluorine.
2F₂ (g) + 2 H₂O (l) → 4H⁺ (aq) + 4 F (aq) + O₂ (g)
X₂ (g) + H₂O (l) → HX (aq) + HOX (aq) (where X = Cl or Br)
4I^– (aq) + 4H⁺ (aq) + O₂ (g) → 2I₂ (s) + 2 H₂O (l)

Anomalous behaviour of Fluorine: It is due to

1. Highest electronegativity and Ionisation enthalpy in Group 17.
2. Small size.
3. Non-availability of d-orbitals.
4. The low bond dissociation energy of the F-F bond.

1. Reactivity towards hydrogen: They all react with hydrogen to give hydrogen halides, but an affinity for hydrogen decreases from F to I.

The acidic strength is in order HF < HCl < HBr < HI and stability is under HF > HCl > HBr > HI.

2. Reactivity towards oxygen: Halogens form many oxides with oxygen but most of them are unstable. F forms OF₂ and O₂F₂, but only OF₂ is stable. Oxides of Cl, Br, I am powerful oxidizing agents.

3. Reactivity towards metals
Mg + Br₂ (l) → MgBr₂(s)
The ionic character of halids decreases in the order MF > MCI > MBr > MI.

4. Formation of interhalogen compounds: They form interhalogen compounds among themselves of the type AB, AB₃, AB₅, AB₇ where A = larger size halogen and B = smaller size halogen.

Chlorine, Cl₂:
Preparation:
1. By heating MnO₂ with cone. HCl

However a mixture of common salt and cone. H₂SO₄ is used in place of HCl.
2 NaCl + MnO₂ + 3 H₂SO₄ → 2 NaHSO₄ + MnSO₄ + 2H₂O + Cl₂

2. By the action of HC1 on KMn04
2 KMnO₄ + 16HCl → 2 KCl + 2 MnCl₂ + 5 Cl₂ + 8 H₂O

Manufacture of Cl₂:
1. Deacon’s process:

2. Cl₂ is obtained by the electrolysis of blue solution.

Properties:
1. Greenish-yellow gas with a pungent smell.

2. 2.5 times heavier than air.

3. Can be liquified easily into greenish-yellow liquid.

4. Soluble in water.

5. It reacts with a number of metals and non-metals and other compounds.
2Al + 3Cl₂ → 2Al Cl₃
2 Na + Cl₂ → 2NaCl
2 Fe +3 Cl₂ → 2FeCl₃
P₄ +6 Cl₂ → 4PCl
S₃ +4 Cl₂ → 4 S₂Cl₂
H₂ + Cl₂ → 2HCl
H₂S + Cl₂ → 2HCl + S
C₁₀H₁₆ + 8 Cl₂ → 16HCl + 10 C.
8NH₃ (excess) + Cl₂ → 6NH₄Cl + N₂
NH₂  + 3Cl₂ (excess) → NCl₃ + 3HCl

6. Chlorine dissolved in water on standing loses its yellow colour due to the formation of HCl and HOCl. Hypochlorous acid (HOCl) so formed, gives nascent oxygen which is responsible for oxidising and bleaching properties: Its oxidising properties are:

(a) 2 FeSO4 + H₂SO₄ + Cl₂ → Fe₂(SO₄)₃ + 2 HCl
Na₂SO₃ + Cl₂ + H₂O → 4 Na₂SO₄ + 2 HCl
SO₂ + 2H₂O + Cl₂ → 4 H₂SO₄ + 2HCl
I₂ + 6 H₂O + 5 Cl₂ → 2 HIO₃ + 10 HCl.

(b) It is a powerful bleaching agent due to oxidation. It is permanent.

Uses:

1. For bleaching wood pulp, cotton and textiles.
2. In the extraction of gold and platinum.
3. Manufacture of dyes, drugs, DDT, refrigerants etc.
4. In sterilising water.

Hydrogen Chloride, HCl:
Preparation-Lab.
Method:
2 NaCl + H₂SO₂  → Na₂SO₄ + 2 HCl cone.

Properties:

1. It is a colourless gas with a pungent smell.
2. Highly soluble in water.
3. Can be liquified easily to a colourless liquid and solidified to a white crystalline solid.
4. It ionises in an aqueous solution.
   HCl + H₂O (l) → 4 H₃O+ (aq) + Cl^– (aq); Ka = 107
   The high value of Ka shows it is a strong acid in water.
5. Reaction with NH₃
   NH₃ + HCl → NH₄Cl (White dense fumes)
6. 3 parts of cone. HCl and one part of cone. HNO₃ on mixing give aqua regia which is used to dissolve, Au, Pt.
   Au + 4H⁺ + NO₃ + 4 Cl^– → AUCl₄^– + NO + 2H₂O
   3Pt + 16H⁺ + 4NO₃^– + 18Cl^– → 3PtCl₆^2- + 4NO + 8H₂O .
7. Reaction with salts:
   Na₂CO₃ + 2HCl → 2 NaCl + H₂O + CO₂ ↑
   NaHCO₃ + HCl → NaCl + H₂O + CO₂ ↑
   Na₂SO₃ + 2HCl → 2 NaCl + H₂O + SO₂ ↑

Uses of HCl:

• Manufacture of Cl₂, NH₄Cl and glucose from com starch.
• Extracting glue from bones and purifying bone black.
• In medicine and as a lab. reagent

Table: Oxo-Acids of Halogens

Structures of Oxoacids:

Inter Halogen Compounds:

Some Properties of Interhalogen Compounds:

Uses:

• Interhafogen compounds are very useful fluorinating agents.
• They can be used as non-aqueous solvents.
• ClF₃ and BrF₃ are used for the production of UF₆ in the enrichment of ²³⁶U.
  U (s) + 3 ClF₃ (l) → UF₆(g) + 3 Cl F (g) argon (Ar)

Group 18 Elements: Group 18 Elements are helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe) and radon (Rn). Because of ns2nph structure [1s² for He], They are chemically unreactive. They are called noble gases.

→ Occurrence: Except radon (Rn) which is radioactive, all noble gases occur in the atmosphere in traces. Noble gases constitute 1 % (by volume) of dry air. He is present in pitch blende, monazite, cleveite. Radon is obtained as a decay product of ²²⁶Ra.
²²⁶₈₈Ra → ²²²₈₆Ra + ⁴₂He

→ Electronic Configuration: All noble gases have ns²np⁶ electronic configuration except He which is 1s².

→ Ionisation Enthalpy: They have very high I.E. which decreases down the group.

→ Atomic radii: They increase down the group with an increase in at. number.

→ Electron Gain Enthalpy: They have large positive values because of their stable electronic configuration. They have no tendency to accept electrons.

Physical Properties:

1. These gases are monoatomic.
2. They are colourless, odourless and tasteless.
3. They are sparingly soluble in water.
4. Due to weak dispersion forces, they have very low m.pts and b. pts.
5. He has the lowest B.pt (4.2 K) of any known substance.

→ Chemical Properties: In general, noble gases are least reactive. It is due to

1. Stable electronic configuration of ns2np6 (or 1s² for He).
2. They have high Ionisation enthalpy and more positive electron gain enthalpy,

Neil Bartlett prepared the first compound of Xe: red coloured Xe⁺ Pt F₆^– by mixing PtF₆^– and Xenon. The compounds of Kr are fewer. He, Ne and Ar are not found to form any true compounds. Only KrF₂ has been studied. Compound involving Rn viz. RnF₂ has been identified but not isolated.

Xenon-fluorine Compounds: Xenon forms three binary compounds with fluorine: XeF₂, XeF₄, XeF₆ by the direct combination of elements under appropriate experimental conditions.

Xenon-oxygen compounds: Xenon and oxygen form XeO₃. Partial hydrolysis of XeF₄ and XeF₆ gives oxyfluorides XeOF₄ and XeO₂F₂.


structures of. (a) XeF₂, (b) XeF₄, (c) XeF₆, (d) XeOF₄ and (e) XeO₃
XeO₃ is a colourless explosive solid and has a pyramidal molecular structure. XeOF₄ is a colourless volatile liquid and has a square pyramidal molecular structure.

Uses: He is used in filling balloons because He is a light and non-inflammable gas. It is used in gas-cooled nuclear reactors. Liquid He is used in low temp, physics. It is used in the driving apparatus.

Neon is used in discharge tubes and fluorescent bulbs for display purposes. Argon is mainly used to provide an inert atmosphere in high-temperature metallurgical processes.

There are no significant uses of Xenon and Krypton. They are used in light bulbs designed for special purposes.

By going through these CBSE Class 12 Chemistry Notes Chapter 6 General Principles and Processes of Isolation of Elements, students can recall all the concepts quickly.

General Principles and Processes of Isolation of Elements Notes Class 12 Chemistry Chapter 6

The extraction and isolation of an element from its combined form involves various principles of chemistry. Metallurgy is the scientific and technological process of extracting a metal from its ore. The natural materials in which the metals or their compounds occur in the earth are called minerals. The mineral from which the metal is extracted conveniently and economically is called an ore. Ores are usually contaminated with earthly or undesired materials known as Gangue.

The extraction and isolation of metals from their respective ores involve the following major steps :

1. The concentration of the ore
2. Isolation of the metal from its concentrated ore, and
3. Purification of the metal.

Principle Ores of Some Important Metals:

Among metals, aluminium is the most abundant. For the purpose of extraction, bauxite is chosen for aluminium. For iron, usually oxide ores [Haematite Fe203] are taken. Before proceeding for concentration, ores are graded and crushed to a reasonable size.

Concentration Of Ores: Removal of the unwanted materials (e.g. sand, clays etc.) from the ore is called concentration, dressing, or benefaction. Unwanted impurities present in the ore are called gangue.

The nature of the impurities, the type of the metal and the environmental factors are taken into consideration.
(a) Hydraulic Washing: it is based on the difference in gravities of the ore and the gangue particles. It is, therefore, a type of gravity separation. When a stream of water is run through the powdered ore, the lighter gangue particles are washed away and heavier ore particles settle down.

(b) Magnetic Separation: This procedure is based on the difference in the magnetic properties of the ore and impurities present in it. When passed over the conveyer belt of a magnetic roller, magnetic particles settle in a heap nearer and non-magnetic impurities a bit away, as shown below:

Magnetic Separation

Froth Floatation Process: This process is used for removing gangue from sulphide ores only. Powdered sulphide ore is mixed with water to which Collectors (e.g. pine oil, fatty acid, xanthates etc.) enhance the non-wettability of the mineral particles and froth stabilizers (e.g. cresols, aniline) which stabilize the froth ore added. The mineral particles become wet by oils while the gangue particles by water.

A rotating paddle agitates the mixture and draws air in it. As a result, froth is formed which carries the mineral particles. The froth is light and is skimmed off from where ore particles are recovered after drying it

Froth Flotation Process

Leaching: This process consists of treating the powdered ore with a suitable reagent (such as acids bases or other chemicals) which can n selectively dissolve the ore but not the impurities.

In Baeyer’s process, pure aluminium oxide is obtained from the bauxite ore (which contains impurities of Fe2Os and silicates) by treating the powdered ore with a 45% solution of NaOH when alumina dissolves leaving behind impurities like Fe203 which are filtered off.

Na [Al(OH)4] or NaAlO., is neutralized by CO2 when Al(OH)3 gets precipitated.

Al(OH)3, obtained above is filtered, washed and finally heated to about 1473 K to get pure alumina Al₂O₃

Extraction of Crude Metal From Concentrated Ore: The concentrated ore must be converted into a form that is suitable for a reduction. Sulphide ore is usually converted into its oxide before reduction. Oxides are easier to reduce.

Two steps are involved
(a) Conversion of the concentrated ore into oxide ore
(b) Reduction of the metal oxide into metal

(a) Conversion To Oxide:
1. Calcination: It is the process of converting ore into its oxide by heating it strongly below its melting point either in the absence or limited supply of air. Volatile matter is driven off.

2. Roasting: It is the process of converting ore into its metallic oxide by heating strongly at a temperature insufficient to melt in excess of air.
2 ZnS + 3O₂ → 2 ZnO + 2 SO₂
2 PbS + 3O₂ → 2 PbO + 2SO₂
2 Cu₂S + 3O₂ → 2 Cu₂O + 2 SO₂

(b) Reduction of oxide to the metal: It is done with a suitable reducing agent (C or CO or even another metal)
M_xO + yC → xM + yCO

Thermodynamic Principles Of Metallurgy: The change in Gibbs energy ΔG is given by
ΔG = ΔH – TΔS …(1)
where ΔH = Enthalpy change at temperature T
ΔS = Entropy change at temperature T

Also ΔG° = -RT ln K. ……….(2)
where R = Gas constant;
K = Equilibrium constant.
The reaction will proceed if ΔG° is negative K in that case will be positive.

This happens only when the reaction proceeds towards products.
1. When the value of ΔG° is negative in equation (1) then only the reaction will proceed. If ΔS is positive, on increasing the temperature (T), the value of TΔS would increase (ΔH < TΔS) and then ΔG will become -ve.

2. If reactants and products of two reactions are put together in a system and the net ΔG of the possible reactions is – ve, the overall reaction will occur. So the process of interpretation involves the coupling of the two reactions, getting the sum of their ΔG and looking for its magnitude and sign. Such coupling is easily understood through Gibbs energy (ΔG°) vs – T plots for the formation of the oxides.

H.J. T Ellingham gave a graphical representation of Gibbs energy. It provides a sound basis for considering the choice of reducing agent in the reduction of oxides. Such a diagram helps us in predicting the feasibility of the thermal reduction of ore.

The reducing agent forms its oxide when the metal oxide is reduced. The role of the reducing agent is to provide ΔG° negative and large enough to make the sum of ΔG° of the two reactions (oxidation of the reducing agent and reduction of the metal oxide) negative.

As we know, during reduction, the oxide of a metal decomposes.
M_xO (s) → xM (Solid or Liquid) + \(\frac{1}{2}\) O₂(g) ……..(3)

The reducing agent takes away the oxygen.
xM (s or l) + \(\frac{1}{2}\) O₂ → MxO(s) ……[ΔG°(M, MxO) ………(4)

If reduction is carried out through equa lion (3), the oxidadion of the reducing agent (e.g. C or CO) will be as:
C(s) + \(\frac{1}{2}\) Oz(g) → CO(g) …………[ΔG(C, CO)]…..(5)
CO(g) + \(\frac{1}{2}\) Oz(g) → CO₂(g)……[ΔG(C, CO)]…(6)

If carbon is taken, there may be complete oxidation to CO2
\(\frac{1}{2}\) C(s)+ \(\frac{1}{2}\)O₂ (g) → \(\frac{1}{2}\)CO₂(g) ….[2 ΔG(C, CO)]…(7)

On subtracting equation (4), we get
M_xO(s) + C(s) → xM(s or l) +CO(g) …(8)
M_xO (s) + C (s) → xM (s or l) + CO₂ (g) …(9)
M_xO (s) + \(\frac{1}{2}\)C(s) → xM (s or l) + \(\frac{1}{2}\)CO₂ (g). …..(10)

These reactions describe the actual reduction of the metal oxide M_xO that is to be accomplished.
Increasing T (Heating) favours a negative value of Δ_rG°. Therefore, the temperature is chosen such that the sum of Δ_rG° in the two combined redox process is negative.

Applications:
(a) Extraction of Iron from its Oxides: Concentrated oxide ores of iron are mixed with limestone and coke and fed into a Blast furnace from the top. Here coke reduces oxide to the metal as follows:
FeO(s) + C(s) → Fe(s or l) + CO(g) ……….(11)

In two steps:
1. FeO(s) → Fe(s) + \(\frac{1}{2}\)O₂ (g) [ΔG(FeO, Fe)] …….(2)
2. C(s) + \(\frac{1}{2}\) O₂ (g) → CO(g) [ΔG(C, CO)] ……..(3)

From (12) and (13), The ne.t Gibbs energy change becomes
ΔG(C, CO) + ΔG(FeO, Fe) = Δ_rG ……….(14)

∴ The resultant reaction will take place when the r.h.s. of equation (14) becomes negative.

The reactions occurring in the Blast furnace at different temperatures are as follows:
At 500 – 800 K (lower temp, range)
3 Fe₂O₃ + CO → 2 Fe₃O₄ + CO₂
Fe₃O₄ + 4CO → 3 Fe + 4 CO₂
Fe₂O₃ + C0 → 2Fe + CO₂

At 900 -1500 K (higher temp, in the blast furnace)
C + CO₂ → 2CO
FeO + CO → Fe + CO₂

Limestone is also decomposed to CaO which removes silicate impurity of the ore as slag. The slag is in a molten state and separates out from iron.
CaCO₃ (s) → CaO (s) + CO₂ (g)
CaO (s) + SiO₂ (s) → CaSiO₂  (fusible slag)

Iron obtained from Blast furnace contains 4% carbon and many impurities in smaller amount (e.g. S, P, Si, Mn) is called pig iron. Cast iron is different from pig iron. It has a slightly lower carbon content (3%) and is extremely hard and brittle.

Blast Furnace

Further Reductions: Wrought iron or malleable iron is the purest form of commercial iron and is prepared from cast iron by oxidising

impurities in a reverberatory furnace lined with haematite. This haematite oxidises carbon to carbon monoxide :
Fe₂O₃ + 3 C → 2 Fe + 3CO

Limestone is added as a flux and sulphur, silicon and phosphorus are oxidised and passed into the slag. The metal is removed and freed from the slag by passing through rollers.

(b) Extraction of Copper from Cuprous Oxide [Copper (I) Oxide]: Most of the copper ores are sulphide ores, concentrated sulphide ores are roasted/smelted to give oxides.
2 Cu₂S + 3 O₂ → 2 Cu₂O + S

The oxide can then be easily reduced to give Cu metal.
Cu₂0 + C_coke → 2 Cu + CO

In the actual process, the ore is heated in a reverberatory furnace after mixing with silica.
In the furnace, iron oxide slags off’ as iron silicate and copper is produced in the form of copper matte which contains Cu₂S and FeS.
FeO + SiO₂ → Fe Si O₃(slag)

In the silica-lined convertor, copper matte is charged. Some silica is added and a hot air blast is blown when the following reactions take place.
2 FeS + 3O₂ → 2 FeO + 2SO₂
FeO + SiO₂ → FeSiO₃
2 Cu₂S + 3O₂ → 2 Cu₂O + 2 SO₂
2 Cu₂O + Cu₂S → 6 Cu + SO₂
The solidified copper obtained has a blistered appearance due to the evolution of SO₂ and so it is called blister copper.

(C) Extraction of Zinc from Zinc Oxide: Zinc oxide is reduced using coke

The metal is distilled off and collected by rapid chilling.

Electrochemical Principles Of Metallurgy: In electrolysis, metal ions in solution or molten form are reduced or by adding some reducing element. Here
ΔG° =. – n FE° ….(i)
where n = no. of electrons transferred
E° = Standard e.m.f. of the cell [redox couple]
F = Faraday = 96,500 C
ΔG° = Change in standard Gibbs energy.

More reactive metals have large negative values of the electrode potential. So their reduction is difficult. If the difference of two E° values corresponds to a positive E° and consequently negative ΔG° in the equation
1. above, then the less reactive metal will come out of the solution and more reactive metal will go into the solution.
Cu²⁺ (aq) + Fe (S) → Cu (S) + Fe²⁺ (aq)

Examples:
Metallurgy of Aluminium: Purified Al₂O₃ is mixed with Na₃ AlF₆ or CaF₂ to, lower the melting point of the mix and to make the solution conductive. The fused matrix is electrolysed.
2 Al₂O₃ + 3C → 4 Al + 3 CO₂

This process of electrolysis is widely known as the Hall-Fieroult process.

Steel cathode and graphite anode are used. The graphite anode is useful here for the reduction of oxide to the metal.

Electrolytic cell for the extraction of aluminium

The electrolysis of the molten mass is carried out in an electrolytic cell using carbon electrodes. The oxygen liberated at the anode reacts with the carbon of the anode producing CO and CO₂.This way for each kg of aluminium – produced, about 0.5 kg of carbon anode is burnt away.

The electrolytic reactions are:
Cathode: Al³⁺ (melt) + 3e^– → Al (l)
Anode: C(s) + O^2- (melt) → CO(g) + 2e^–
C(s) + 2 O^2- (melt) → CO₂ (g) + 4e^–

Copper from low-grade ores and scraps: Copper is extracted from low grades ores by Hydrometallurgy. It is leached out using acid or bacteria. The solution containing Cu²⁺ (aq) is treated with iron or H₂.
Cu²⁺ (aq) + H₂ (g) → Cu (s) + 2H⁺(aq)

Oxidation-Reduction: Some non-metals are extracted based on oxidation. Chlorine from prime solution is oxidized to Cl₂.
2 Cl^– (aq) + 2 H₂O (l) → 2 OH^– (aq) + H₂ (g) + Cl₂ (g)

In the extraction of gold and silver, metal is leached with CN^–.
Ag → Ag⁺ + e^– oxidation
Au → Au⁺ + e^– oxidation

The metal is later recovered by the displacement method.
4 Au (s) + 8 CN^– (aq) + 2H₂O (aq) + O₂ → 4 [Au (CN)₂]^– (aq) + 4 OH^– (aq)
2 [Au (CN)₂]^– (aq) + Zn (s) → 2 Au (s) + [Zn (CN)₄]^2- (aq)
In this reaction zinc acts as a reducing agent.

Refining: To obtain metal of high purity and to remove the last traces of impurities from the extracted metal, they are subjected to refining.

It is based upon the difference in properties of the metal and the impurity.
Several techniques are listed below:

1. Distillation,
2. Liquation,
3. Electrolysis,
4. Zone refining,
5. Vapour phase refining,
6. Chromatographic methods.

1. Distillation: This is useful for low boiling metals like Zn and Hg.

2. Liquation: A low melting metal like tin (Sn) can be made to flow on a sloping surface leaving higher melting impurities.

3. Electrolytic Refining: Impure metal is made anode. A strip of pure metal is made the cathode. The electrolyte (bath) contains soluble salt of the same metal.
Anode: M → Mn⁺ + e^–
Cathode: Mn⁺ + ne^– → M

Copper is refined by the electrolytic method. Here acidified solution of copper sulphate acts as an electrolyte.
Anode: Cu(s) → Cu²⁺ + 2e
Cathode: Cu²⁺ + 2e → Cu(s)

4. Zone Refining: This method is based upon the principle that the impurities are more soluble in the melt than in the solid state of the metal. The molten zone moves along with the mobile heater fixed at one end of the impure metal. As the heater moves forward the pure metal crystallises out of the metal and the impurities pass on into the adjacent molten zone. At one end, impurities get concentrated. This end is cut off.

Zone-Refining Process

Vapour Phase Refining: Mond’s Process for Nickel Refining

The carbonyl complex is subjected to a higher temperature so that it is decomposed to give pure metal.

van Arkel Method for Refining Zirconium or Titanium: This method is very useful for removing all the oxygen and nitrogen present in the form of impurity in certain metals like Zr and Ti. The crude metal is heated in an evacuated vessel with iodine. The metal iodide, being more covalent, volatilises.
Zr + 2I₂ → ZrI₄.

Chromatographic Methods: This method is based on the principle that different components of a mixture are differently adsorbed on an adsorbent.

Column Chromatography: It is very useful for the purification of the elements which are available in minute quantities and the impurities are not very different in chemical properties from the element to be purified.

There are several chromatographic techniques such as

1. Paper chromatography
2. Gas chromatography

Uses of Aluminium, Copper, Zinc and Iron: Alloys containing Al are light and are very useful. A1 wires conduct electricity. It is used as a reductant.

Copper is used to making wires in the electrical industry. Several alloys of copper with Zn, Sn and Ni are largely used.
Zinc is used for galvanising iron. It is used as a reducing agent. Similarly different forms of iron: Cast iron, Wrought iron, Steel find wide applications.

A Summary of the Occurrence and Extraction of Some Metals is Presented in the following Table:

By going through these CBSE Class 12 Chemistry Notes Chapter 5 Surface Chemistry, students can recall all the concepts quickly.

Surface Chemistry Notes Class 12 Chemistry Chapter 5

The branch of chemistry which deals with the nature of surface and species present on it is called surface chemistry. Adsorption on solid or on solution surfaces and colloidal properties are important surface effects.

Adsorption: The phenomenon of higher concentration of molecular species (gases or liquids) on the surface of solids than in the bulk is called adsorption.

The solid on the surface of which adsorption occurs is called adsorbent. The substances that get adsorbed on the solid surface is called adsorbate. The adsorbent may be a solid or a liquid and the adsorbate may be a gas or a liquid.

The phenomenon of absorption differs from adsorption as:

Absorption	Adsorption
1. It is the phenomenon in which the particles of gas or liquid get uniformly distributed throughout the body of the solid.	1. It is the phenomenon of higher concentration of particles of gas or liquid on the surface than if! the bulk of the solid.
2. Absorption occurs at uniform rate.	2. Adsorption is rapid in the beginning and its rate slowly decreases.

Types Of Adsorption:
Depending upon the nature of forces between the molecules of the adsorbate and the adsorbent, the adsorption maybe classified as: physical adsorption and chemical adsorption.
1. Physical adsorption: When the particles of the adsorbate are held to the surface of the adsorbent by the weak forces such as van der Waals forces, the adsorption is called physical adsorption or physisorption. The attractive forces are weak and therefore, these can be easily overcome either by increasing the temperature or by decreasing the pressure. In other words, physical adsorption can be easily reversed or decreased.

2. Chemical adsorption: When the molecules of the adsorbate are held to the surface of the adsorbent by the chemical forces, the adsorption is called chemical adsorption or chemisorption, hi this case, a chemical reaction occurs between the adsorbed molecules and the molecules or atoms of adsorbent.

Table: Comparison between Physisorption and Chemisorption:

Physisorption	Chemisorption
1. Enthalpy of adsorption, usually is of the order of 20 – 40 kJ mol-1	1. Enthalpy of adsorption, is of the order 40-400 kJ mol-1.
2. Molecules of adsorbate and adsorbent are held by van der Waals interactions.	2. Molecules of adsorbate and adsorbent are held by chemical bonds.
3. It usually takes place at low temperature and decreases with increasing temperature.	3. It takes place at relatively high temperatures.
4. It is not very specific i.e., all gases are absorbed on all solids to some extent.	4. It is highly specific and takes place when there is some possibility of compound formation between the adsorbate and the adsorbent molecules.
5. Multi-molecular layers may be formed on the adsorbent.	5. Usually mono-molecular layer is formed on the adsorbqnt,
6. It does not require any activation energy.	6. It requires activation energy.
7. The amount of gas adsorbed is related to the ease of liquefaction of the gas.	7. There is no such correlation.
8. it is reversible in nature.	8. It is irreversible in nature.

Freundlich Adsorption Isotherm:
The adsorption of a gas on the surface of the solid depends upon the pressure of the gas. The extent of adsorption is generally expressed as x/m where m is the mass of the adsorbent and x is the mass of adsorbate when equilibrium has been attained. On the basis of experimental studies, Freundlich gave the following relationship between the amount of gas adsorbed (x) per unit mass of the adsorbent (m) and the pressure (p).
\(\frac{x}{m}\) = kp^1/n
where n is a constant (whole number) which depends upon the nature of adsorba te and adsorbent.

A graph between the amount (x/m) adsorbed by an adsorbent and the equilibrium pressure (or concentration for solutions) of the adsorbate at constant temperature is called Adsorption Isotherm.

At low pressure \(\frac{x}{m}\) ∝ p¹ ………..(i)

At high pressure \(\frac{x}{m}\) ∝ p⁰ ………..(ii)

In the intermediate range of pressure, combining (i) and (ii)
\(\frac{x}{m}\) ∝ p^0-1
∝ p1/n where n is an integer (n > 1)
or
\(\frac{x}{m}\) = kp^1/n ………..(iii)
where k is a constant depending upon the nature of the adsorbate and adsorbent.

This relationship is called Freundlich Adsorption Isotherm
Taking logs on both sides of (iii)
log \(\frac{x}{m}\) = log k + \(\frac{1}{n}\) log p

A graph between log \(\frac{x}{m}\) against log p should, therefore, be a straight line with slope equal to \(\frac{1}{n}\) and ordinate intercept equal to log K.

Adsorption From Solution Phase:
Solids can adsorb solutes from solutions so when a solution of acetic acid in water is shaken with charcoal, a part of the acid is adsorbed by charcoal and the concentration of the acid decreases in the solution.

The following conclusions have been made regarding adsorption from solution phase:

1. The extent of adsorption decreases with increase in temperature.
2. The extent of adsorption increases with the increase in the surface area of the absorbent.
3. The extent of adsorption depends upon the concentration of the solute in the solution.
4. The extent of adsorption depends upon the nature of the absorbent and the adsorbate.

Freundlich equation as applied to solutions is modified as \(\frac{x}{m}\) = k C^1/n  i.e., where C is the equilibrium concentration,
or
log \(\frac{x}{m}\) = log k + \(\frac{1}{n}\) log C

Applications of Adsorption:
1. Production of high vacuum: The remaining traces of air can be adsorbed by charcoal from a vessel evacuated by a vacuum pump to give a very high vacuum.

2. Gas masks: Gas mask (a device which consists of activated charcoal or mixture of adsorbents) is usually used for breathing in coal mines to adsorb poisonous gases.

3. Control of humidity: Silica and aluminium gels are used as adsorbents for removing moisture and controlling humidity.

4. Removal of colouring matter from solutions: Animal charcoal removes colours of solutions by adsorbing coloured impurities.

5. Heterogeneous catalysis: Adsorption of reactants on the solid surface of the catalysts increases the rate of reaction. There are many gaseous reactions of industrial importance involving solid catalysts. Manufacture of ammonia using iron as a catalyst, manufacture of H₂SO₄ by contact process and use of finely divided nickel in the hydrogenation of oils are excellent examples of heterogeneous catalysis.

6. Separation of inert gases: Due to the difference in degree of adsorption of gases by charcoal, a mixture of noble gases can be separated by adsorption on coconut charcoal at different temperatures.

7. In curing diseases: A number of drugs are used to kill germs by getting adsorbed on them.

8. Froth floatation process: A low grade sulphide ore is concentrated by separating it from silica and other earthy matter by this method using pine oil and frothing agent (see Unit 6).

9. Adsorption indicators: Surfaces of certain precipitates such as silver halides have the property of adsorbing some dyes like eosfn, fluorescein, etc. and thereby producing a characteristic colour at the end point.

10. Chromatographic analysis: Chromatographic analysis based on the phenomenon of adsorption finds a number of applications in analytical and industrial fields.

Catalysis: Potassium chlorate when heated strongly decomposes slowly giving dioxygen. The decomposition between 653 – 873 K.
2 KClO₃ → 2 KCl + 3O₂

However with a little of MnO a decomposition occurs only at 473 – 633 K and also at a much accelerated rate. MnO2 is a catalyst for this reaction.

A substance which accelerates the rate of a chemical reaction, itself remaining chemically and quantitatively unchanged after the reaction is called a catalyst.

Promoter is a substance that enhances the activity of a catalyst, while Poison is a substance that decreases the activity of a catalyst. In the reaction for the manufacture of NH₃ by Haber’s process.

Fe is catalyst and Molybdenum (MO) is a promoter.

Homogeneous And Heterogeneous Catalysis:
When the reactants and the catalyst are in the same phase (liquid or gas), the process is said to be homogeneous catalysis.

Example:

When the reactants and the catalyst are in different phases, the process is called heterogeneous catalysis.

Important Features of solid catalysts are:
(a) Activity: It depends upon the strength of chemisorption to a large extent. .
(b) Selectivity: Selectivity of a catalyst is its ability to direct a reaction to yield a particular product. For example, starting with H₂ and CO and using different catalysts, we get different products.

Thus a catalyst is highly selective in nature, i.e., a given substance can act as a catalyst only in a particular reaction and not for all the reactions. A catalyst for a particular reaction may fail to catalyse any other reaction.

Shape-Selective Catalysis by Zeolites: The catalytical reaction that depends upon the pore structure of the catalyst and the size of the reactant and product molecules is called shape selective catalysis. Zeolites are good shape-selective catalysts because of their honey comb¬like structure.

An important Zeolite catalyst used in the petroleum industry is ZSM-5. It converts alcohols directly into gasoline (petrol) by dehydrating them to give a mixture of hydrocarbons.

Enzyme Catalysis: Enzymes are complex nitrogeneous organic compounds which are produced by living plants and animals. They are actually protein molecules of high molecular mass, and form colloidal solutions in water. The enzymes are also referred to as Bio-Chemical Catalysts as they also occur in the bodies of animals and plants and such a phenomenon is known as Bio-Chemical Catalysis.

The following are examples of enzyme-catalysed reactions:
1. Inversion of cane-sugar

2. Conversion of glucose into ethyl alcohol.

3. Decomposition of urea into ammonia and CO₂.

Enzyme	Source	Enzymatic Reaction
Invertase	Yeast	Sucrose → Glucose and fructose
Zymase	Yeast	Glucose → Ethyl alcohol and carbondioxide
Diastase	Malt	Starch → Maltose
Maltase	Yeast	Maltose → Glucose
Urease	Soyabean	Urea → Ammonia and carbon dioxide
Pepsin	Stomach	Proteins → Amino acids

Characteristics Of Enzyme Catalysis:
Enzyme catalysis is unique in its efficiency and high degree of specificity. The following characteristics are exhibited by enzyme catalysts:
1. Most highly efficient: One molecule of an enzyme may transform one million molecules of the reactant per minute.

2. Highly specific nature: Each enzyme is specific for a given reaction, i.e., one catalyst cannot catalyse more than one reaction. For example, the enzyme urease catalyses the hydrolysis of urea only. It does not catalyse hydrolysis of any other amide.

3. Highly active under optimum pH: The rate of an enzyme- catalysed reaction is maximum at a particular pH called optimum pH, which is between pH value 5-7.

4. Highly active under optimum temperature: The rate of an enzyme reaction becomes,maximum at a definite temperature, called the optimum temperature. On either side of the optimum temperature, the enzyme activity decreases. The optimum temperature range for enzymatic activity is 298-310 K. Human body temperature being 310 K is suited to enzyme-catalysed reaction.

5. Increasing activity in presence of activators and co¬enzymes: The enzymatic activity is increased in the presence of certain substances, known as co-enzymes. It has been observed that when a small non-protein (vitamin) is present along with an enzyme, the catalytic activity is enhanced considerably.

Activators are generally metal ions such as Na⁺ Mn²⁺, CO²⁺, Cu²⁺, etc. These metal ions, when weakly bonded to enzyme molecules, increase their catalytic activity. Amylase in presence of sodium chloride i.e., Na⁺ ions are catalytically very active.

6. Influence of inhibitors and poisons: Like ordinary catalysts enzymes are also inhibited or poisoned by the presence of certain substances. The inhibitors or poisons interact with the active functional groups on the enzyme surface and often reduce or completely destroy the catalytic activity of the enzymes. The use of many drugs is related to their action as enzyme inhibitors in the body.

Mechanism of Enzyme Catalysis: There are a number of cavities present on the surface of collodial particles of enzymes. These cavities are of characteristic shape and possess active groups such as – NH₂, – COOH, – SH, – OH etc. These are actually the active centres on the surface of enzyme particles. The molecules of the reactant (substrate), which have complementary shape, fit into these cavities just like a key fits into a lock.

An activated complex is formed which then decomposes to yield the products in two steps as outlined below:
Step I: E + S → ES*
Step II: ES* → E + P

Some Industrial Catalytical Processes:

Process	Catalyst
1. Haber’s process for the manu-facture of ammonia N2(g) + 3H2(g) → 2NH3(g)	1. Finely divided iron, molybdenum as promoter; conditions: 200 bar pressure and 723 – 773K temperature.
2. Ostwald’s process for the manu-facture of nitric acid. 4NH3(g) + 5O2(g) → 4N0(g) + 6H2O(g) 2NO(g) + O2(g) → 2NO2(g) 4NO2(g) + 2H2O(l) + O2(g) → 4HNO3(l)	2. Platinised asbestos; temperature: 573K.
3. Contact process for the manu-facture of sulphuric acid. 2SO2(g) + O2 (g) ⇌ 2SO3(g) SO3(g) + H2SO4 (l) → H2S2O7 (l)oleum H2S2O7 (l)+H2O (l) → 2H2SO4 (l)	3. Platinised asbestos or vanadium pentoxide (V205); temperature 673-723K.

Colloids: A colloid is a heterogeneous system in which one substance is dispersed (dispersed phase) as very fine particles in another substance called dispersion medium.

The essential difference between a solution and a colloid is that of particle size. Their size is in between that of true solution and suspension

Classification of Colloids: Colloids are classified on the basis of the following criteria:
(a) Physical state of dispersed phase and dispersion medium.
(b) Nature of interaction between dispersed phase and dispersion medium.
(c) Type of particles of the dispersed phase.
(a) Physical state of dispersed phase and dispersion medium: Depending upon whether the dispersed phase and the dispersion medium are solids, liquids or gases, eight types of colloidal systems are possible.

Types of Collodial Systems:

(b) Depending upon the nature of interactions between the dispersed phase and dispersion medium, the colloids can be classified as Lyophilic Colloids and Lyophobic Colloids.

1. Lyophilic collids: The colloidal solutions in which the particles of the dispersed phase have a great affinity (or love) for the dispersion medium are called lyophilic colloids. Such solutions are reversible in nature. In case water acts as the dispersion medium, the lyophilic colloid is called hydrophilic colloid. The common examples of lyophilic colloids are glue, gelatin, starch, proteins, rubber etc.

2. Lyophobic colloids: The colloidal solutions in which the particles of the dispersed phase have no affinity or love, rather have hatred for the dispersion medium, are called lyophobic collids. The solutions of metals like Ag and Au, hydroxides like Al (OH)₃ and Fe (OH)₃ and metal sulphides like As₂S₃ are examples of lyophobic colloids.
Such sols are formed with difficulty. They are irreversible in nature.

Multimolecular Macromolecular And Associated Colloids:
Depending upon the molecular size, the colloids can be classified as:
1. Multimolecular colloids: In this type, the particles consist of an aggregate of atoms or small molecules with molecular size less than 1 nm. For example, sols of gold atoms and sulphur (S₈) molecules. In these colloids, the particles are held together by Van der Waals forces.

2. Macromolecular colloids: In this type, the particles of the dispersed phase are sufficiently big in size (macro) to be of colloidal dimensions. In this case, a large number of small molecules are joined together through their primary valencies to form giant molecules.

These molecules are called macro molecules and each macromolecule may consist of hundreds or thousands of simple molecules. The solution of such moleucles are called macromolecular soluions. For example, colloidal solution of starch, cellulose, etc.

3. Associated colloids: These are the substances which behave as normal electrolytes at low concentration but behave, as colloidal particles at higher concentration. These associated particles are also called miscelles. For example, in aqueous solution, soap (sodium stearate) ionises as:

In concentrated solutions, these ions get associated to form an aggregate of colloidal size.

General Methods Of Preparation Of Sols:
Lyophilic sols are readily formed by simply mixing the dispersed phase and the dispersion medium under ordinary conditions.

Lyophobic sols can generally be prepared by two methods.
1. Condensation methods,
2. Dispersion methods.

→ Condensation methods: In these methods, the smaller particles are condensed suitably to be of colloidal size. This can be done by chemical reactions or by exchange of solvent.

→ Dispersion methods: In these methods, the large particles of a substance (suspension) are broken into smaller particles. This can be done by mechanical dispersion, by electrical dispersion or Bredig’s arc method and by peptisation.

Purification Of Colloidal Solutions:
The colloidal solutions prepared usually contain impurities especially electrolytes which can destabilize the sols. These impurities must be eliminated to make the colloidal solution stable.

The following methods are commonly used for the purification of colloidal solutions.
1. Dialysis: The method is based upon the fact that colloidal particles cannot pass through a parchment or cellophane membrane while the ions of the electrolyte can pass through it. The colloidal solution is taken in a bag made of cellophane or parchment.

The bag is suspended in fresh water. The impurities slowly diffuse out of the bag leaving behind pure colloidal solution. For example, dialysis can be used for removing HCl from the ferric hydroxide sol.

2. Electrodialysis: The ordinary process of dialysis is slow:

An apparatus for electrodialysis

To speed up the process of purification, the dialysis is carried out by applying electric field. This process is called electrodialysis.

3. Ultra-filtration: It is the process of removing the impurities from the colloidal solution by passing it through graded filter paper called ultrafilter papers. These filter papers are made from ordinary filter papers by impregnating them with colloidal solutions.

As a result, the size of the pores gets reduced. These filter papers allow the ions and molecules of the impurities to pass but retain colloidal particles. Ordinary filter papers cannot be used for this purpose since the colloidal particles also easily pass through the pores of these papers.

Properties Of Colloidal Solutions: The important properties of colloidal solutions are:
1. Heterogeneous nature: The colloidal solutions are heterogenous in nature consisting of dispersed phase and dispersion medium.

2. Visibility: The colloidals are not visible to naked eye and these can be seen with ultra microscopes.

3. Brownian movement: The colloidal particles have continuous zigzag motion called Brownian movement.

Brownian Movement

4. Tyndall effect: When the light is passed through the colloidal solution, the path of the light becomes visible when viewed from a direction at right angle to the incident beam. This phenomenon was studied by Tyndall and is known as Tyndall effect. The phenomenon of scattering oflight’by colloidal particles as a result of which the path of the beam becomes visible is called Tyndall effect.

5. Electrical properties: The particles of the colloidal solutions possess electrical charge, positive or negative. The presence of charge is responsible for the stability of these solutions. It may be noted that only the sol particles carry some charge while the dispersion medium has no charge. For example, the collodial solutions of gold, arsenious sulphide (AS₂S₃) are negatively charged while those of Fe (OH)₃ and Al (OH)₃ have positive charge. In the case of silver chloride sol, the particles may either be positively or negatively charged.

The presence of the charge on the sol particles and its nature whether positive of negative can be determined with the help of a phenomenon known as electrophoresis. In this experiment, the colloidal particles move towards positive or negative electrodes depending upon their charge under the influence of electrical field.

The phenomenon of movement of colloidal particles under an applied electric field is called electrophoresis.

If the particles accumulate near the negative electrode, the charge on the particles is positive. On the other hand, if the sol particles accumulate near the positive electrode the charge on the particles is negative.

A set up for electrophoresis

Origin of charge: The charge on the colloidal particles may be due to selective adsorption of ions. The particles contributing the dispersed phase adsorb only those ions preferentially which are common with their own lattice ions. For example, if silver nitrate solution is added to an aqueous solution of potassium iodide, the silver iodide will adsorb negative ions (I^–) from the dispersion medium to form a negatively charged sol.

However, if silver iodide is formed by adding potassium iodide to silver nitrate solution, the sol will be positively charged due to the adsorption of Ag⁺ ions present in the dispersion medium.

Coagulation Of Colloidal Solution: The phenomenon of precipitation of a colloidal solution by the addition of excess of an electrolyte is called coagulation or floculation.

→ Factors governing coagulation
1. Nature of the electrolytes: The coagulation capacity of different electrolytes depends upon the valency cf the active ion or called floculating ion. It is the ion carrying charge opposite to the charge on the colloidal particles. According to Hardy Schulz law, greater the valency of the active ion or floculating ion greater will be its coagulating power. Thus, to coagulate negative sol of As₂S₃, the coagulating power of different cations has been found to decrease in the order as:
Al³⁺ > Mg²⁺ > Na⁺

Similarly, to coagulate a positive sol such as Fe (OH)₃, the coagulating power of different anions has been found to decrease in the order:
[Fe(CN)₆]^4- > PO₄ ^3- > SO₄^2- > Cl^–

The minimum concentration of an electrolyte which is required to cause the coagulation or flocculation of a sol is known as flocculation value. It is usually expressed as milli moles per litre.

Protection of Colloids: The process of protecting the lyophobic colloidal solution from precipitation by the electrolytes due to the previous addition of some lyophilic colloid is called protection. The colloid which is added to achieve such a protection is called protecting colloid.

Gold Number: The different protecting colloids differ in their portecting powers. Zsigmondy introduced a term called gold number to describe the protective power of different colloids. This is defined as the minimum number of milligrams of the protective colloid required to just prevent the coagulation of a 10 ml of a given gold sol when 1 ml of a 10% solution of sodium chloride is added to it.

The coagulation of gold sol is indicated by change incolour from red to blue. The gold number of a few protective colloids are as follows:

It may be noted that smaller the value of gold number, greater will be protecting power of the protective colloid. Therefore, reciprocal of gold number is a measure of the protective power of a colloid. Thus, out of the list given above, gelatin is the best protective colloid.

Emulsions: Emulsions are the colloidal solutions of two immiscible liquids in which the liquid acts as the dispersed phase as well as the dispersion medium. Normally they are obtained by mixing an oil with water. Since the two do not mix well, the emulsion is generally unstable and is stabilised by adding a suitable outside reagent called emulsifier or emulsifying agent. The substances that are commonly employed for the purposes are gum, soap, glass powder, etc.

Types of Emulsions: These are of two types:
1. Oil-in-water emulsions: In this case, oils acts as the dispersed phase (small amount) and water as the dispersion medium (excess) e.g., milk is an emulsion of soluble fats in water and here casein acts as an emulsifier. Vanishing cream is another example of this class. Such emulsions are called aqueous emulsions.

2. Water-in-oil emulsions: In this case water acts as the dispersed phase while the oil behaves as the dispersion medium e.g., butter, cod liver oil, cold cream etc. Such types of emulsions are called oily emulsions.

Types of Emulsions

Demulsification: It is the process of decomposing an emulsion back into its constituent liquids. The demulsification can be done by centrifugation, filtration, boiling, freezing and some chemical methods.

Following are the interesting noteworthy examples of colloids which we come across in daily life.

1. Blue colour of the sky: Dust particles along with water suspended scatter blue light due to which sky looks blue to, us.
2. Fog, mist and rain: Clouds are aerosols. It is possible to cause artificial rain by throwing electrified sand.
3. Food articles like milk, butter, fruit juices are all colloids.
4. Blood-: Blood is a colloidal solution of an albuminoid substahce. Alum and FeCl₃ solution stop oozing blood due to coagulation. ,
5. Soils: Fertile soils are colloidal in nature in which humus . acts as a protective colloid.
6. Formation of delta: River water is a colloidal solution of clay.

Sea water contains several electrolytes. When river water meets sea water, the electrolytes present in sea water, coagulate the colloidal solution of clay resulting in its deposition with the formation of delta.

Applications of colloids: Colloids are widely used in the industry.
Some examples are:
1. Cottrell Smoke Precipitator: Smoke is a colloidal solution of solid particles such as carbon, arsenic compounds, dust etc. in air. The smoke is led through a chamber containing platgs having a charge opposite to that carried by smoke particles.” The particles on coming in contact with these plates lose their charge and get precipitated. The particles thus settle down on the floor of the chamber.

Cottrell Smoke Precipitator

2. Purification of drinking water: Alum is added to water to coagulate the suspended impurities present in water from natural resources to make it fit for drinking purposes.

3. Medicines: Most of the medicines are colloidal in nature. . Milk of magnesia, an emulsion, is used for stomach disorder. Argyrol is a silver sol used as an eye lotion. Colloidal antimony is used in curing Kalazar. Colloidal gold is used for intramuscular injection. Colloidal medicines are more effective because of large surface area and hence are easily assimilated.

4. Tanning: Animal hides due to colloidal nature bear positive charge when it is soaked in tanning or chromium salts which Contain negatively charged colloidal particles, mutual coagulation takes place leading to hardening of leather. The process is called tanning.

5. Cleansing action of soaps and detergents: Soaps and detergents act as emulsifiers and remove greasy impurities and dust of clothes which is washed away.

6. Photographic plates & films are prepared by coating an emulsion of light sensitive AgBr in gelatin over glass plates or celluloid films.

7. Rubber industry: Latex is a colloidal solution of rubber particles which are negatively charged. Rubber is ob tained by coagulation of latex.

8. Industrial products: Paints, inks, synthetic plastics and rubber, graphite lubricants, cement etc. are all colloidal solutions.

By going through these CBSE Class 12 Chemistry Notes Chapter 4 Chemical Kinetics, students can recall all the concepts quickly.

Chemical Kinetics Notes Class 12 Chemistry Chapter 4

Thermodynamics tells us about the feasibility of a reaction. A reaction is feasible if ΔG < 0 at constant temperature and pressure. Chemical equilibrium tells us the extent to which a reaction can proceed. But none of them tell us the speed of a reaction, i.e., time taken by a reaction to reach equilibrium.

The branch of chemistry which deals with the rates of reactions and their mechanisms is called Chemical Kinetics.
Rate Of A Chemical Reaction

From a kinetics point of view, reactions can be classified into 3 categories:
1. Very fast reactions: Which take place instantaneously e.g., ionic reactions like
AgNO₃ (aq) + NaCl (aq) → ↓AgCl (s) + NaNO₃ (aq)
BaCl₂ (aq) + H₂SO₄ (aq) → BaS04 (s) + 2 HCl (aq)
NaOH (aq) + HCl (aq) → NaCl (aq) + H₂O (l)

2. Very slow reactions: Which may take days or months for

Decomposition of NH₄NO₂

Decomposition of Hydrogen peroxide

Rate Of A Chemical Reaction: The rate of the reaction means the speed of the reaction.
It is defined as the change in the concentration of any one of the reactants or products per unit time.
Rate of the reaction = \(\frac{\text { Decrease in the concentration of a reactant }}{\text { Time taken }}\)
= \(\frac{\text { Increase in the concentration of a product }}{\text { Time taken }}\)

Consider a hypothetical reaction, R → P assuming the volume of the system is constant. One mole of reactant R produces one mole of the product P. If [R]₁ and [P]₁ are the concentrations of R and P at time t₁ and [R]₂ and [P]₂ are their concentrations at time t₂ then
ΔT = t₂ – t₁
Δ[R] = [R]₂ – [R]₁
Δ[P] = [P]₂ – [P]₁
The square brackets are used to express molar concentrations.
Rate of disappearance of R = – \(\frac{\Delta[\mathrm{R}]}{\Delta \mathrm{t}}\) …….(i)
Rate of appearance of P = + \(\frac{\Delta[\mathrm{P}]}{\Delta \mathrm{t}}\) ………(ii)

Equations (i) and (ii) given above represent the average rate of a. reaction, r_ave. r_ave depends upon (a) change in the cone, of reactants or products and (ii) time taken for that change to occur.

Units of Rate of a Reaction: Units of rate are concentration time-1, i.e., mol L^-1 S^-1. However for gaseous reactions, when the concentrations are expressed in terms of partial pressure, then the units of the rate will be atm s^-1.

For example,
PCl₅ → PCl₃ + Cl₂

Any general reaction
A + B → C + D

For this reaction, the rate of decomposition (disappearance) of

Rate of formation of O₂ = + \(\frac{\mathrm{d}\left[\mathrm{NO}_{2}\right]}{\mathrm{d} \mathrm{t}}\)

Rate law: The dependence of the concentration of reactants on the rate of reaction is given by the law. It gives the experimental dependence on concentration of reactants. For example, for the reaction
aA + bB → Products

From the kinetic study of the reaction, the dependence of the concentration of reactants on the rate of the reaction has been found to be

Rate = k[A]^m[B]ⁿ
where m and n are constant numbers or the powers of the concentration terms of the reactants A and B respectively on which the rate of reaction depends. It may be noted that the values of m and n are determined experimentally and may or may not be equal to a and b coefficients in the reaction. The above expression is the rate law. It may be defined as:

The mathematical expression denotes the observed or actual rate of a reaction in terms of the molar concentration of the reacting species which influences the rate of the reaction.

Order And Molecularity Of A Reaction:
Order of a Reaction: The dependence of reaction rate on concentration may be expressed in terms of the order of a reaction. The order of a reaction is defined as

the sum of the powers to which the concentration terms in the rate law are raised to express the observed rate of a reaction.
Thus, if the rate of a reaction,
xA + yB + zC → Products is given by the rate law,

Rate = –\(\frac{\mathrm{d} \mathrm{x}}{\mathrm{dt}}\) = k[A]x [B]y [C]z then, the order of the reaction, n is
n = x + y + z
where x, y and z are the order with respect to individual reactants and overall order of the reaction is sum of these exponents, i.e., x + y + z.

For example,
1. decomposition of ammonium nitrite is a first-order reaction.
NH₄NO₂ → N₂ + 2H₂O
Rate = k [NH₄NO₂] Order = 1

2. decomposition of hydrogen iodide is a second order reaction
2HI → H₂ + I₂
Rate = k[HI]₂ Order = 2

3. the reaction between nitric oxide and oxygen is a third-order reaction.
2NO (g) + O₂ (g) → 2NO₂ (g)
Rate = k[NO]²[O₂] Order = 3

The order of a reaction may be fractional or even zero.
For example,
CH₃CHO → CH₄ + CO
Rate = k[CH₃CHO]^3/2 Order = 3/2

Rate = k [NH₃]° Order = 0

Molecularity of Reaction: The molecularity of a reaction is defined as the number of reacting molecules that collide simultaneously to bring about a chemical reaction.

For example, the decomposition of H₂O₂ involves only one molecule, so it is a unimolecular reaction.
H₂O₂ → H₂O + \(\frac{1}{2}\)O₂

Dissociation of hydrogen iodide is a biomolecular reaction.
2HI (g) → H₂ (g) + I₂ (g)

The reaction involving three or more than three molecules is uncommon because the occurrence of such reactions would require simultaneous collisions of three or more than three molecules. The chances of the occurrence of such collisions are very small.

However, some of these reactions are found to be quite fast. This means that even though the balanced equation involves a large number of molecules, yet the reaction does not proceed by the simultaneous collision of all these reacting particles.

Such types of reactions take place through a sequence of two or more consecutive steps and are called complex reactions. The detailed description of various steps by which reactants change into the products is called the mechanism of the reaction. The steps which contribute to the overall reaction are called elementary processes.

Mechanism and Rate Law: In the case of multi-step reactions, some of the steps will be very fast while others will be slow. If one step takes place much more slowly than all other steps, it controls the overall reaction rate. This means that all the steps have to wait for the occurrence of this slowest step. But once this slowest step has occurred, the other steps will take place to form the products. In other words, the rate of the reaction is determined by the slowest step in the sequence.

Let us consider reaction between NO₂ and F₂ to form NO₂F;
2NO₂(g) + F₂(g) → 2NO₂F(g)

The experimental observations reveal that the rate of the reaction is proportional to the product of the concentrations of nitrogen peroxide and fluorine. This indicates that the rate-determining step in the mechanism of this reaction must be the reaction between NO₂ and F₂ only. Keeping this in mind, a mechanism of this reaction may be suggested as:

The rate of the overall reaction is determined by the first step which is the slower of the two steps. Accordingly, the experimentally observed rate of the reaction is given by the expression :
fix
Rate = – \(\frac{\mathrm{d} x}{\mathrm{dt}}\) = k[NO₂][F₂]

The above is the rate law for the reaction. It is, therefore, evident that complex reactions involving more than three molecules in the stoichiometric equation must take place in more than one step.
e.g.KClO₃ + 6 FeSO₄ + 3 H₂SO₄ → KCl + 3Fe₂ (SO₄)₃ + 3H₂O
should be, apparently, of the tenth order is actually a second-order reaction. This shows that the reaction occurs in several steps. The slowest step is called the rate-determining step.

Consider

The rate equation for this reaction is found to be
Rate = – \(\frac{\mathrm{d}\left[\mathrm{H}_{2} \mathrm{O}_{2}\right]}{\mathrm{dt}}\) = k [H202][I-]

This reaction is of first order w.r.t. both H2Oz and I-. Evidences suggest that this reaction takes place in 2 steps :

1. H₂O₂ + I → H₂O + IO^– – Slow Step
2. H₂O₂ + IO^– → H₂O + I^– + O₂ – Fast Step

∴ The first step, being slow, is the rate determining step.

It can be concluded from above

1. The order of a reaction is an experimental quantity. It can be zero and even a fraction, but molecularity cannot be zero or a non-integer.
2. Order is applicable to elementary as well as complex reactions whereas molecularity is applicable only for elementary reactions. For complex reactions, molecularity has no meaning/significance.
3. For a complex reaction, the order is given by the slowest step and generally, the molecularity of the slowest step is the same as the order of the overall reaction.

Units of Reaction Rate Constants: The rate is the change in concentration with time. Therefore, the rate of reaction is expressed by concentration units divided by time. If the concentrations are expressed in mol litre^-1 and time in seconds, then units of rate constants for different orders are:

• Units for first-order reaction = sec^-1
• Units for second-order reaction = litre mol^-1 sec^-1
• Units for third-order reaction = litre² mol^-2 sec^-1

In the case of gaseous reactions, if concentrations are expressed in units of atm, then

• Units of rate constant for the first-order reaction = sec^-1
• Units of rate constant for second-order reaction = atm-1 sec^-1
• Units of rate constant for third-order reaction = atm^-2 sec^-1

Rate Constant For First Order Reaction:
For a first-order reaction,
A → Products
Let the initial cone, of [A] be an ML^-1.

Let x change into products so that equilibrium concentration after time t is


If the initial concentration is ‘a’ moles per liter, x moles of A change in time t and k is the rate constant, then the integrated rate equation is

where [A]₀ is the initial concentration and [A] is the concentration at time t. The value of k can be calculated by substituting the values of a, t, and x.

A plot between In (A) and t gives a straight line with a slope equal to – k.

Half-Life Period Of A Reaction: The half-life period of a reaction is defined as the time during which the concentration of a reactant is reduced to one-half of its initial concentration. It is generally denoted as t_1/2. The half-life period of a first-order reaction may be calculated as given below:

For the first-order reaction,
t = \(\frac{2.303}{\mathrm{k}}\) log \(\frac{[\mathrm{A}]_{0}}{[\mathrm{~A}]}\)

Now half-life period corresponds to the time during which the initial concentration, [A]₀ = a, is reduced to half i.e. [A] = a/2

The half-life period, t_1/2becomes

Thus, half life period of a first-order reaction is independent of the initial concentration of the reactant.

Similarly, the relation for the time required to reduce the concentration of the reactant to any fraction of the initial concentration can be calculated. For example,
∴ t_3/4 = \(\frac{2.303}{\mathrm{k}}\) log \(\frac{\mathrm{a}}{\mathrm{c} / 4}\) = \(\frac{2.303}{\mathrm{k}}\) log 4

Zero order Reactions are those reactions in which the rate of the reaction is proportional to zero power of the concentration of the reactants.
Rate = – \(\frac{\mathrm{d}[\mathrm{R}]}{\mathrm{dt}}\) = k[R]_o = k
or
R = \(\frac{[\mathrm{R}]_{0}-[\mathrm{R}]}{\mathrm{t}}\)

Collision Theory of Reaction Rate: The number of collisions that take place per second per unit volume of the reaction mixture is known as collision frequency (Z). All the collisions are not effective and do not give products. The collisions which actually produce the products and therefore, result in chemical reactions are called effective collisions.

There are two conditions for effective collisions.
1. Energy barrier: For the reacting species to make effective collisions, they should have sufficient energy to break the chemical bond in the reacting molecules. The minimum amount of energy that the colliding molecules must possess is known as threshold energy. Thus, only those collisions of reactants will give products that possess energies greater than threshold energy.

2. Orientation barrier: The colliding molecules should also have proper orientation so that the old bonds may break and new bonds are formed.

Thus, the collisions in which the colliding molecules do not possess the threshold energy of proper orientation do not form products. Therefore, only a small fraction of collision is effective.

Orientation of colliding NO₂ molecules is proper in (a) but ‘not’ in (b) for the reaction to take place.

Dependance Of Reaction Rate On Temperature: Temperature has a great influence on reaction rates. In general, the rate of a reaction becomes almost double for every 10° rise in temperature. The increase in reaction rate is not due to an increase in collision frequency but it is due to an increase in the fraction of effective collisions. It has been found that the fraction of effective collisions becomes almost double for a 10° rise in temperature.

Arrhenius Equation: The quantitative relationship between the rate constant and temperature was proposed by Arrhenius known as the Arrhenius equation.
k = A e^-E_a/RT

where k is a constant called frequency factor, E_a is the activation energy. If k₁ and k₂ are rate constants at two different temperatures T₁ and T₂ respectively then, integrated form of Arrhenius equation is
log\(\frac{\mathrm{k}_{2}}{\mathrm{k}_{1}}=\frac{\mathrm{E}_{\mathrm{a}}}{2.303 \mathrm{R}}\left[\frac{\mathrm{T}_{2}-\mathrm{T}_{1}}{\mathrm{~T}_{1} \mathrm{~T}_{2}}\right]\)

Activation Energy: The excess of energy (over and above the average energy of the reactants) required by the reactants to undergo chemical reactions is called activation energy. It is equal to the difference between threshold energy needed for the reaction and the average energy of reactant molecules.

Activation Energy = Threshold energy – Average energy of all reacting molecules.

When the molecules possess the energy equal to E , the atomic configuration of species formed at this stage is different from the reactants as well as the products. This stage is called the activated state or the transition state and specific configuration of the state is called Activated Complex. For example in the reaction between H₂ (g) and I₂ (g), activated complex has configuration in which H-H arid 1-1 bonds are breacking and H-I bonds are forming as shown below.

The change of reactants to products, i.e., progress of a reaction is shown below.

E_a = Activation energy for forward reaction
E_a‘ = Activation energy for backward reaction

Effect Of Catalyst On Reaction Rate:
The substances which increase the rate of a reaction and can be recovered chemically unchanged in mass and composition after the reaction are called catalysts. The phenomenon of increasing the rate of reactions by the use of catalyst is known as catalysis.

The function of a catalyst is that it provides a new path for the reaction, in which the reactants are converted into products quickly. It is believed that the catalyst forms a new activated complex of lower potential energy. This means that the activation energy becomes lower for the catalysed reaction than that for uncatalysed reaction.

Consequently, the fraction of the total number of collisions possessing lower activation energy is increased and hence, the rate of reaction also increases. This is shown below. The solid lines shows the path for uncatalysed reaction and dotted line shows the path adopted by catalysed reaction.

E_a = Activation energy without catalyst
E’_a (Cat) = Activation energy with catalyst

Effect Of Radiation On Reaction Rate:
Rate of reaction is increased by use of certain radiations. Such type of reactions which are initiated by the absorption of radiation are called photochemical reactions. For example, reaction of hydrogen and chlorine takes pla,ce very slowly in the absence of light. However, in the presence of light, the reaction occurs rapidly.

Reversible Reactions: At any stage, the net rate of the reaction is determined by the difference in the rates of forward and backward reactions. At equilibrium the overall rate of reaction in either direction becomes zero. Therefore,
Net rate of reaction = Rate of forward reaction – Rate of backward reaction At equilibrium, net rate is zero, so that

Rate of forward reaction = Rate of backward reaction
For the simple reaction A + B ⇌ C + D

If the rate constant for the forward reaction is Rf and that for the backward reaction is kf then.
Rate of forward reaction a [A] [B] = k [A] [B]
Similarly, rate of backward reaction a [C] [D] = kb [C] [D]
∴ At equilibrium rate of forward reaction = rate of backward reaction

where K = Equilibrium constant
Thus, K = \(\frac{[C] \times[D]}{[A] \times[B]}\)

By going through these CBSE Class 12 Chemistry Notes Chapter 3 Electrochemistry, students can recall all the concepts quickly.

Electrochemistry Notes Class 12 Chemistry Chapter 3

Electrochemical Cell: It is a device to convert the chemical energy of a spontaneous redox reaction into electrical energy. It is also called Galvanic Or Voltaic Cell. One such example of an electrochemical cell is Daniel Cell.

Daniel’s cell converts the chemical energy liberated during the redox reactions.
Zn (s) + Cu²⁺ (aq) → Zn²⁺ (aq) + Cu (s) into electrical energy and has an electrical potential [E.M.F. ] equal to 1.1 volt when concentration of both Zn²⁺ and Cu²⁺ ions is unity (1 mol dm^-3)

In the left-hand beaker, oxidation occurs. Zn plate dissolves to form Zn²⁺. Zn plate loses weight
Oxidation Half-reaction:
At anode: Zn (s) → Zn²⁺ (aq) + 2e^–

Reduction Half-Reaction: In the right-hand beaker reduction occurs. Cu²⁺ ions from the solution deposit on the copper plate. It gains weight.
At Cathode: Cu²⁺ (aq) + 2e^– → Cu (s)

Electrons flow from zinc plate to copper plate in the external circuit. Conventional current flows from copper to zinc plate (as. shown above) Combining the two half-reactions, we get the complete cell reaction or redox reaction.
Zn (s) + Cu²⁺ (aq) → Zn²⁺ (aq) + 2e^–

Electrolytic Cell: It is a device to convert electrical energy from an external source to produce a chemical change. In the above cell if E_ext is > 1.1. the reaction will again start but in the opposite reaction. It now will function as an electrolytic cell.

(a) when Eext < 1.1 V

1. Electrons flow from Zn rod to Cu rod hence current flows from Cu to Zn.
2. Zn dissolves at anode and copper deposits at the cathode.

(b) when E_ext = 1.1 V

1. No flow of electrons or current.
2. No chemical reaction.

Functioning of Daniell cell when external voltage E_ext opposing the cell potential is applied.
(c) When E_ext > 1.1 V

1. Electrons flow from Cu to Zn arid current flows from Zn to Cu.
2. Zinc is deposited at the zinc electrode and copper dissolves at the copper electrode

In the electrochemical cell at each electrode-electrolyte interface, there is a tendency of metal ions from the solution to deposit on the metal electrode trying to make it positively charged. At the same time, metal atoms of the electrode have a tendency to go into the solution as ions and leave behind the electrons at the electrode trying to make it negatively charged.

At equilibrium, there is a separation of charges and depending on the tendencies of the two opposing reactions, the electrode may be positively or negatively charged with respect to the solution. A potential difference develops between the electrode and the electrolyte which is called Electrode Potential.

When the concentration of all species involved in half-cells is unity, the electrode potential is called Standard Electrode Potential.

According to the IUPAC convention, standard reduction potentials are now called standard electrode potentials. In a galvanic cell, the half¬cell in which oxidation takes place is called Anode and it has a negative potential with respect to the solution. The other half-cell in which reduction takes place is called Cathode and it has a positive potential with respect to the solution. The direction of the flow of current is opposite to that of the flow of electrons.

The cell potential is the difference between the electrode potentials (reduction potentials) of the cathode and anode. It is called electromotive force (emf) of the cell when no current flows through the circuit. Internally the two half-cells/beakers are connected through the salt bridge.
E_cell = E_right – E_left

In the Cu^– AgNO₃ cell cell reaction is
Cu (s) + 2 Ag⁺ (aq) → Cu²⁺ (aq) + 2 Ag (s)
At anode: Cu (s) → Cu²⁺ (aq) + 2 e^–] oxidation
At cathode: 2 Ag⁺ (aq) + 2 e^– → 2Ag (s)] reduction

Here Ag electrode acts as cathode and copper electrode, as anode.
The cell can be represented by
Cu (s) | Cu²⁺ (aq) || Ag⁺ (aq) | Ag (s)
E_cell = E_right – E_left
= E_Ag⁺/Ag – E_Cu²⁺/Cu

Similarly Daxtiel cell can be represented by
Zn (s) | Zn²⁺ (aq) || Cu²⁺ (aq)| Cu (s)
E_cell = E_right – Eleft
= E_Cu²⁺/Cu – E_Zn²⁺/Zn
Measurement of Electrode Potential: The measurement of the potential of half-cells is not possible.

Only the difference between the two half-cell potentials can be determined that gives the e.m.f. of the cell. By arbitrarily fixing the electrode potential of one half-cell, that of the other can be determined.

According to the convention, Standard Hydrogen Electrode represented by Pt (s), H₂(g), H₂(g)/H⁺ (aq) is assigned a zero potential at all temperatures according to the reaction
H⁺(aq) + e^– → H₂(g)

Standard Hydrogen Electrode: It consists of platinum wire sealed in a glass tube arid has a platinum foil attached to it. The foil is coated with finely divided platinum and acts as a platinum electrode. It is dipped into an acid solution containing H⁺ ions in 1 M concentration (1M HCl). Pure H₂ gas at 1 bar pressure is constantly bubbled into the solution at a constant temperature of 298 K. The following reaction occurs in the half-cell depending upon whether it acts as an anode or as a cathode.

Normal Hydrogen Electrode (NHE)

If S.H.E. (orN.H.E.) acts as Anode
H₂(g) → 2H⁺ (aq) + 2e^– oxidation half reaction

If S.H.E. (or N..H.E.) acts as Cathode
2H⁺ (aq) + 2e^– → H₂ (g)] reduction half reaction

The standard hydrogen electrode is also regarded as a reversible electrode with respect to H+ ions:
H₂(g) ⇌ 2H⁺ (aq) + 2e^–

Arbitrarily, the standard electrode potential of this electrode is fixed to be 0.000 V.
The electrode potential of an electrode can be determined by connecting this half-cell with a standard hydrogen electrode.

The electrode potential of a metal electrode as determined with respect to S.H.E. (or N.H.E.) is called Standard Electrode Potential E°.

Conventionally: The reduction (standard) potential of an electrode that acts as a Cathode when attached to S.H.E. is given a positive sign, e.g., Cu/Cu⁺⁺ electrode when attached to S.H.E. acts as a cathode.

Let us calculate the electrode (reduction) potential of Zn/Zn²⁺. The cell will be
Zn (s) | Zn²⁺_1.0 (aq)| | H⁺ _1.0 (aq), H₂ (g), Pt(s).

Let us measure the electrode potential of Cu/Cu²⁺ electrode concentration of Cu²⁺ is 1.0 M and the pressure of H₂ gas is one bar. At 298 K, the emf of the cell

Standard hydrogen electrode 11 second half cell.
E°_cell = E°_R – E°_L
0.34 = E°_Cu⁺⁺ – 0 [emf of the cell = 0.34 V]
∴ E_Red° of Cu electrode dipping in Cu²⁺ ions = 0.34 V

Electrochemical Series or The Standard Electrode Potentials at 298 K: Ions are present as aqueous species and H₂O as liquid, gases and solids are shown by aq, g and s.

1. A negative Ev means that the redox couple is a stronger reducing agent than the H⁺/ H₂ couple.
2. A positive Ev means that the redox couple is a weaker reducing agent than the H⁺/H₂ couple.

Applications Of Electrochemical Series:
1. Relative strength of oxidizing and reducing agents: In this series, metals are arranged in the decreasing order of their standard electrode potentials or decreasing order of their oxidizing character. e.g. Li⁺ (aq) which has least reduction potential (- 3.05) is the weakest oxidizing agent and F₂ (g) which has maximum value of p_Red = + 2.85 is the strongest oxidizing agent.

On the other hand, Li metal which has the highest standard oxidation potential (= + 3.05) is the strongest reducing agent while F- ( E°_ox = – 2.85 V) is the least reducing agent.

2. Calculation of e.m.f. of the cell: The e.m.f. of the desired cell can be calculated knowing the standard reducing potentials of the two half cells constituting the cell from the electrochemical series.
E°_cell = E°_Red [RIGHT] – E°Red [LEFT]
e.g. Zn (s) | Zn²⁺ (1.0 M) || Cu²⁺ (1.0 M) | Cu (s).
E°_cell = E°_Cu²⁺,_Cu – E°_Zn²⁺,_Zn.
= 0.34 – (- 0.76) V = 1.10 V

3. Predicting feasibility of redox reaction: In general, a redox reaction is feasible only if the species which has higher standard reduction potential is reduced i.e., accepts the electrons and the species which has lower reduction potential is oxidized i.e., loses the electrons. Otherwise, a redox reaction is not feasible. In other words, the species to release the electrons must have lower reduction potentials as compared to the species which is to accept electrons. In a nutshell, if the e.m.f. of the hypothetical cell is +ve, a redox reaction takes place; if it is – ve, a redox reaction is not feasible.

4. To predict whether a metal can librate H₂(g) from acid or not. The metals which have only negative reduction potentials, i.e., are lying above N.H.E. in the electrochemical series, can only liberate H₂ (g) from dilute acid solutions.

Nernst Equation: The relationship between electrode potentials and the concentration of the electrolytic solutions is called Nernst Equation.
1. The Nernst Equation is
E = E° + \(\frac{\mathrm{RT}}{\mathrm{nF}}\) In [Mn+]/[M] Since[M] = 1
E = E° + \(\frac{2.303 \mathrm{RT}}{\mathrm{nF}}\) log [Mn+]
E = reduction electrode potential
E° = reduction electrode potential in standard state [1 M solution of metal ions at 298 K]
R = Gas constant = 8.314 JK-1.M0l-1
T = Temperature
n = no. of electrons accepted during the change
| F | = Faraday t = 96500 C.

Putting the values of R, T = 298 K; F
| E | = E° + \(\frac{0.059}{\mathrm{n}}\) log [Mn+]

For the complete cell reaction
Zn (s) + Cu²⁺(aq) ⇌ Zn²⁺ (aq) + Cu (s)

2. Calculation of Gibbs free energy/Maximum work that can be obtained from the Galvanic cell
ΛG° = – n F E°elI = – 2.303 RT log Kc
where AG° is the standard Gibbs energy change in Gibbs function
n = number of moles of electrons involved in the cell reaction
F = Faraday = 96,500 coulombs
R = Gas Constant
Kc = Equilibrium constant.

Conductance of Electrolytic Solutions:
Resistance R (ohm) = \(\frac{\text { Potential Difference in volts }}{\text { Current strength in amperes }}\)

Resistance can be measured by the principle of the Wheatstone bridge. It depends upon the length and inversely on the area of cross-section A of the wire.
R ∝ l
R ∝ \(\frac{l}{\mathrm{~A}}\)
or
R = ρ\(\frac{l}{\mathrm{~A}}\);
where ρ (rho) is called specific resistance/Resistivity.

The resistivity of a substance is its resistance to the passage of electricity when it is one metre long and its area of cross-section is one m2.

The inverse of resistance R is called conductance.
Conductance = \(\frac{1}{\mathrm{R}}=\frac{\mathrm{A}}{\rho l}=\frac{\mathrm{A} \kappa}{l}\) ( κ = Kappa)

The SI unit of conductance is Siemens.
1S = 1 ohm^-1 = mho = Ω^-1.

The inverse of resistivity is called Conductivity or Specific Conductance.
SI unit of conductivity is Sm-1, but often (Greek Kappa) is expressed in S cm^-1.

The magnitude of conductivity varies a great deal and depends upon the nature of the material. It also depends upon the temperature and pressure of measurement.

Materials are classified as

1. Conductors,
2. Insulators,
3. Semi-conductors depending upon the magnitude of their conductivity.

Metals and their alloys have very large conductivity and are known as conductors. Certain, non-metals like carbon black, graphite and some organic polymers are also electronically conductors. Substance like glass, ceramics etc. have very low conductivity are known as insulators. Substance like silicon, doped silicon, gallium arsenide is semi-conductors. Electrical conductance ( through metals called metallic or electronic conductance is due to the flow of mobile electrons.

The metallic conductance depends upon,

1. The nature and structure of the metal.
2. The number of valence electrons per atom.
3. Temperature. It decreases with an increase in temperature.
4. The composition of the metallic conductor remains unchanged.

Electrolytic or Ionic Conductance: When electrolytes are dissolved in water, they furnish their own ions in the solution. They conduct electricity through their ions in the solution and is called electrolytic or ionic conductance.

The conductivity of electrolytic or ionic solutions depends upon:

1. The nature of the electrolyte added
2. Size of the ions produced and their solvation.
3. The nature of the solvent and its viscosity.
4. The concentration of the electrolyte.
5. Temperature conductivity increases with an increase in temperature.
6. It leads to a change in its composition with the passage of direct current through the solution over a prolonged period.

Measurement of Conductivity of Ionic Solutions: It is based upon the measurement of resistance by a wheat stone bridge. The cell is called a conductivity cell.

Unknown resistance of the cell R2 = \(\frac{\mathrm{R}_{1} \mathrm{R}_{4}}{\mathrm{R}_{3}}\)
Conductance C of the cell = \(\frac{1}{\mathrm{R}_{2}}\)

Cell Constant = \(\frac{\text { Specific conductance }}{\text { Observated conductance }}\)

Molar Conductivity: It is the conducting power of all ions obtained by dissolving 1 mole of an electrolyte in a given volume of the solution,

Molar conductivity = Λm = \(\frac{x}{c}\)
= \(\frac{\mathrm{Sm}^{-1}}{1000 \mathrm{Lm}^{-3} \times \text { molarity }\left(\mathrm{mL}^{-1}\right)}\)
= S m² mol^-1

Variation of Conductivity and Molar Conductivity with concentration: Both conductivity and molar conductivity change with the concentration of the electrolyte. Conductivity always decreases with the decrease in concentration both for weak and strong electrolytes. This can be explained by the fact that the number of ions per unit volume that carry the current in a solution decreases on dilution. The conductivity of a solution at any given concentration is the conductance of one unit volume of solution kept between two platinum electrodes with a unit area of cross-section and at a distance of unit length. This is clear from the equation:

C = \(\frac{\mathrm{κA}}{l}\)= κ (both A and l are unity in their appropriate Units in m or cm)

Molar conductivity of a solution at a given concentration is the conductance of volume V of a solution containing one mole of electrolyte kept between two electrodes with an area of cross-section A and distance of unit length. Therefore,
Λ_m = \(\frac{\mathrm{κA}}{l}\)= κ
Since l = 1 and A = V (volume containing 1 mole of electrolyte)
Λ_m = κV

Molar conductivity increases with a decrease in concentration. This is because the total volume, V, of a solution containing one mole of electrolyte also increases. It has been found that a decrease in κ on dilution of a solution is more than compensated by an increase in its volume.

Physically, it means that at a given concentration, Λm can be defined as the conductance of the solution of an electrolyte kept between the electrodes of a conductivity cell at a unit distance but having an area of cross-section large enough to accommodate sufficient volume of solution that contains one mole of the electrolyte. When concentration approaches zero, the molar conductivity is known as limiting molar conductivity and is represented by the symbol Λ°_m The variation in Λm with concentration is different for strong and weak electrolytes.

Strong Electrolytes: For strong electrolytes, A increases slowly with dilution and can be represented by the equation :
Λ_m = Λ°_m – A c^1/2
It can be seen that if we plot Λm against c^1/2, we obtain a straight line with an intercept equal to and slope equal to ‘A’. The value of the constant ‘A’ for a given solvent and temperature depends on the type of electrolyte i.e., the charges on the cation and anion produced on the dissociation of the electrolyte in the solution. Thus, NaCl, CaCl₂, MgS04 are known as 1-1, 2-1 and 2-2 electrolytes respectively. All electrolytes Of a particular type have the same value for ‘A’.

Molar conductivity versus c^1/2 for acetic acid (weak electrolyte) and potassium chloride (strong electrolyte) in aqueous solutions.

Kohlrausch law of independent migration of ions. The law states that limiting molar conductivity of an electrolyte can be represented as the sum of the individual contributions of the anion and cation of the electrolyte. Thus, if λ°_Na and λ°_Cl^– are limiting molar conductivity of the sodium and chloride ions respectively, then the limiting molar conductivity for sodium chloride is given by the equation :
Λ°_m(NaCl) = λ°_Na + λ°_Cl^–

In general: Λ°_m = v₊ λ°₊ + n_– λ°_–
Here λ°₊ and λ°_– are the limiting molar conductivities of the cation and anion respectively.

Weak Electrolytes: Weak electrolytes like acetic acid have a lower degree of dissociation at higher concentration and hence for such electrolytes, the change in Λ_m with dilution is due to an increase in the degree of dissociation and consequently the number of ions in the total volume of solution that contains 1 mol of electrolyte. In such cases, Λ_m increases steeply on dilution, especially near lower concentrations.

Therefore, Λ°m cannot be obtained by extrapolation of Λm to zero concentration. At infinite dilution (i.e., concentration c → zero) electrolyte 5 dissociates completely (α = 1), but at such low concentration, the conductivity of the solution is so low that it cannot be measured accurately. Therefore, Λ°_m for weak electrolytes is obtained by using Kohlrausch law of independent migration of ions. At any concentration c, if a is the degree of dissociation then it can be approximated to the ratio of molar conductivity A m at the concentration c to limiting molar conductivity, Λ°_m.

Thus we have:
α = \(\frac{\Lambda_{\mathrm{m}}}{\Lambda_{\mathrm{m}}^{\circ}}\)
But we know that for a weak electrolyte like acetic acid,

Applications of Kohlrausch law: Using Kohlrausch law of independent migration of ions, it is possible to calculate Λ°_m for any electrolyte from the λ° of individual ions. Moreover, for weak electrolytes like acetic acid, it is possible to determine the value of its dissociation constant once we know the Λ°_m and Λ_m at a given concentration c

Electrolytic Cells and Electrolysis: In an electrolytic cell external source of voltage is used to bring about a chemical reaction. The electrochemical processes are of great importance in the laboratory and the chemical industry. One of the simplest electrolytic cells consists of two copper strips dipping in an aqueous solution of copper sulphate. If a DC voltage is applied to the two electrodes, then Cu²⁺ ions discharge at the cathode (negatively charged) and the following reaction takes place:
Cu²⁺ (aq) + 2e^– → Cu (s)
Copper metal is deposited on the cathode. At the anode copper is converted into Cu²⁺ ions by the reaction:
Cu (s) → Cu²⁺ (s) + 2e^–

Thus copper is dissolved (oxidised) at the anode and deposited (reduced) at the cathode. This is the basis for an industrial process in which impure copper is converted into the copper of high purity. The impure copper is made an anode that dissolves on passing current and pure copper is deposited at the cathode.

Many metals like Na, Mg, Al, etc. are produced on large scale by electrochemical reduction of their respective cations where no suitable chemical reducing agents are available for this purpose. Sodium and magnesium metals are produced by the electrolysis of their fused chlorides and aluminium is produced by electrolysis of aluminium oxide in presence of cryolite.

Faraday’s Laws of Electrolysis:
1. First Law: The amount of chemical reaction which occurs at any electrode during electrolysis by a current is proportional to the quantity of electricity passed through the electrolyte (solution or melt).

2. Second Law: The amounts of different substances liberated by the same quantity of electricity passing through the electrolytic solution are proportional to their chemical equivalent weights (Atomic Mass of Metal. Number of electrons required to reduce the cation).

Some Commercial Cells: The electrochemical cells can be used to generate electricity and these are called batteries. The battery is generally used for two or more galvanic cells connected in series.

There are two types of commercial cells:

1. Primary cells, in which electrode reactions cannot be reversed by external energy source. Therefore, these are not chargeable. For example, dry cell, mercury cell.
2. Secondary cells are those which can be recharged. For example, lead storage cell, nickel-cadmium cell.

→ Fuel Cells: These are voltaic cells in which the reactants are continuously supplied to the electrodes. These are designed to convert the energy from the combustion of fuels such as H₂, CO, CH₄ etc. directly into electrical energy. A common example is a hydrogen-oxygen fuel cell.

→ Corrosion: The process of deterioration of a metal as a result of its reaction with air or water surrounding it is called corrosion. In the case of iron, corrosion is called rusting. Chemically rust is a hydrated form of ferric oxide, Fe₂O₃ x H₂O. It is caused by moisture, carbon dioxide and oxygen present in the air.

The chemistry of corrosion is essentially an electrochemical phenomenon. At a particular spot of an iron object, oxidation takes place and that place behaves as an anode:
Anode 2 Fe(s) → 2Fe²⁺ + 4e^– E°_Fe²⁺/Fe = – 0.44 V

Electrons released at anodic spot move through the metal and go to another spot on the metal and reduce oxygen in presence of H⁺ (which is believed to be .available from H2C03 formed due to dissolution of carbon dioxide from the air into water. Hydrogen ion in water may also be available due to the dissolution of other acidic oxides from the atmosphere).

This spot behaves as cathode with the reaction Cathode:
O₂(g) + 4H⁺ (aq) + 4e^– → 2 H₂O
E_H⁺|O2|H2O = 1.23V

The overall reaction being:
2Fe(s) + O₂(g) + 4H⁺ (aq) → 2Fe₂ (aq) + 2 H₂O (l)
E^V_cell = 1.67 V

The ferrous ions are further oxidised by atmospheric oxygen to ferric ions which come out as rust in the form of hydrated ferric oxide (Fe₂O₂ . x H₂0) and with further production of hydrogen ions.

Corrosion of iron in the atmosphere

Oxidation: Fe (s) → Fe²⁺ (aq) + 2e^–
Reduction: O₂(g) + 4H⁺ (aq) + 4e^– → 2H₂O (1)

Atmospheric
oxidation : 2 Fe²⁺ (aq) + 2H₂O (l) + 1/2 O₂ (g) → Fe₂O₃ (s) + 4H⁺ (aq)

Prevention of Corrosion: The rusting of iron can be prevented or decreased by the following methods:
1. Barrier protection: In this method a barrier is placed between the iron and the atmospheric air by coating the surface with paint, applying a thin film of oil or grease or electroplating.

2. Sacrificial protection: In this method, iron is protected by covering it with a layer of metal more active than iron such as zinc, tin, etc. The process of covering iron with zinc is called galvanization.

3. Electrical protection: In this method, the iron is connected with a more active metal like magnesium or zinc.

4. Using anti-rust solution: To retard the corrosion of iron certain anti-rust solutions such as alkaline phosphates and alkaline chromates are used.

Commercial Production of Chemicals: Many metals and their compounds are prepared by using basic principles of electrolysis:

Na is prepared by the electrolysis of molten Na⁺ Cl^– by Down’s Cell

Mg is prepared from fured Mg Cl₂
MgCl₂(Z) ⇌ Mg²⁺ (l) + 2 Cl^– (l)
At anode 2Cl^– → Cl₂ (g) + 2e^–
At cathode Mg²⁺ + 2e^– → Mg (s).

The Hydrogen Economy: Both the production of hydrogen by electrolysis of water
2 H₂O (l) → 2H₂ (g) + O₂ (g) and hydrogen combustion in a fuel cell
2 H₂ (g) + O₂ (g) 2H₂O (l). will be important in future. Both are based on electrochemical principles.

By going through these CBSE Class 12 Chemistry Notes Chapter 2 Solutions, students can recall all the concepts quickly.

Solutions Notes Class 12 Chemistry Chapter 2

Solution: It is a homogeneous mixture of two or more components whose composition may be varied within limits. The component present in the largest quantity is known as Solvent. One or more components present in the solution other than the solvent are called Solutes. Solutions consisting up of two components only are called Binary Solutions.

In all there are 9 types of solutions:
Expressing Concentration of Solutions
1. Mass Percentage (W/W):
Mass of the comp, in solution
Mass % of a component= \(\frac{\text { Mass of the comp. in solution }}{\text { Total mass of solution }}\) × 100

2. Volume Percentage (V/V):
Volume % of a component = \(\frac{\text { Volume of the component }}{\text { Total volume of solution }}\) × 100
The total volume of the solution

3. Mass by volume percentage (W/V):
It is the mass of solute dissolved in 100 mL of the solution.

4. Parts per million (ppm):
Parts per million
\(\frac{\text { No. of parts of the component }}{\text { Total no. of parts of all components of the solution }}\) × 10⁶

5. Mole Fraction:
Mole fraction of a component (x) = \(\frac{\text { No. of moles of the component }}{\text { Total no. of moles of all components }}\)

If there are 2 components in a binary solution of A and B.
x_A = \(\frac{n_{A}}{n_{A}+n_{B}}\)
x_B = \(\frac{n_{B}}{n_{A}+n_{B}}\)
[where nA and nB are the no. of moles of A and B respectively]
x_A + xB_B = 1

For a solution containing i no. of components
x_i = \(\frac{n_{i}}{n_{1}+n_{2}+\ldots n_{i}}=\frac{n_{i}}{\Sigma n_{i}}\)

6. Molarity (M): Molarity is defined as the no. of moles of the solute dissolved in one litre (or one cubic decimetre) of solution.
Molarity (M) = \(\frac{\text { Moles of solute }}{\text { Volume of solution in litre }}\)
= \(\frac{\text { Strength of solute per litre of solution }}{\text { Molar mass of the solute }}\)

7. Molality (m): It is defined as the no. of moles of the solute per 1000 g [l kg] of the solvent.
Molality (m) = \(\frac{\text { Moles of solute }}{\text { Mass of solvent in } \mathrm{kg}}\)
= \(\frac{\text { Strength of the solute per } 1000 \mathrm{~g} \text { of solvent }}{\text { Molar mass of the solute }}\)

Mass %, ppm, mole fraction and molality are independent of temperature, whereas molarity depends upon temperature.

Solubility: The solubility of a substance is its maximum amount that can be dissolved in a given amount of solvent at a given temperature. It depends upon the nature of solute and solvent as well as temperature and pressure.

→ Solubility of a solid in a Liquid: In general, polar solutes dissolve in polar solvents and non-polar solutes in non-polar solvents. Like Dissolves Like. In general, a solute dissolves in a solvent if the intermolecular interactions are similar.
Solute + Solvent → Solution

→ Saturated Solution: A solution is said to be saturated if no more of the solute can be dissolved in it at a particular temperature and pressure.

→ Effect of Temperature on Solubility: The solubility of a solid in a liquid is significantly affected by temperature changes. If the process of dissolution in a nearly saturated solution is endothermic (Δsol H > 0), the solubility should increase with rising temperature and if it is exothermic (Δsol H < 0), the solubility should decrease.

→ Effect of Pressure: Pressure does not have any significant effect on the solubility of solids in liquids. It is so because solids and liquids are highly incompressible and practically remain unaffected by changes in pressure.

→ Solubility of a Gas in a Liquid: The solubility of gases in liquids is greatly affected by pressure and temperature. The solubility of gases increases with the increase of pressure.

Henry’s Law: It states, “At a constant temperature, “The solubility of a gas in a liquid is directly proportional to the pressure of the gas.”

Dalton also concluded, “The solubility of a gas in a liquid solution is a function of the partial pressure of the gas.”
In other words, “The mole fraction of the gas in the solution is proportional to the partial pressure of the gas over the solution.”

The most commonly used form of Henry’s law states “The partial pressure of the gas in the vapour phase (p) is proportional to the mole fraction of the gas (x) in the solution.”
p = K_Hx

Where KH is Henry’s law constant and it is a function of the nature of the gas.

Solubility of a gas in a liquid decreases with an increase in temperature. It is due to this reason that aquatic species like fish are more comfortable in cold waters rather than warm water.

Applications of Henry’s Law:

1. Pressure is kept high to increase the solubility of CO₂ in soft drinks.
2. Deep-sea divers experience bends or decompression sickness due to greater solubility of N₂ and O₂ in blood. O₂ is dissolved in the blood and other body fluids and N₂ will remain dissolved and will form bubbles when the diver comes to the atmospheric pressure. Hence N₂ in compressed air is’ replaced by He which is much less soluble in biological fluids.
3. In the function of lungs: Due to the partial pressure of O₂ being high when it enters the lungs, it combines with haemoglobin to form oxyhaemoglobin. The partial pressure of O₂ in tissues is low. Hence O₂ is released from oxyhaemoglobin which is utilised for functions of cells

Vapour Pressure of Liquid-Liquid Solutions: In a binary solution of two volatile liquids in a closed vessel, let Ptotal, p₁ and p₂ denote total vapour of the solution, partial vapour pressures of the two compounds and x₁ and x₂ their mole fractions, then according to Raoult’s Law.
p₁ ∝ x₁
and p₁ = p₁° × x₁
where p₁° is the V.P. of pure component 1

Similarly p₂ = p₂° × x₂
where p₂° is the V.P. of pure component 2

[Raoult’s Law states “For a solution of volatile liquids, the partial vapour pressure of each component in the solution is directly proportional to its mole fraction.”]

Now according to Dalton’s law of partial pressures.
P_total = P₁ + P₂
= p₁° + p₂° x₂
= (1 – x₂)p° + x₂ p°
= P₁° + ( P₂° – P₁° )x₂.

Following conclusions can be drawn from the above equations

1. Total vapour pressure over the solution can be related to the mole fraction of any one component.
2. Total vapour pressure over the solution varies linearly with the mole fraction of component 2.
3. Depending upon the vapour pressure of the pure components 1 and 2, total vapour pressure over the solution decreases or increases with the increase of the mole fraction of component 1.

Assuming component 1 to be less volatile than component 2, the minimum value of p_total is p° and the maximum value p° (i.e. p° < p° ), a graph can be drawn between the mole fraction and their vapour pressures in solution.

The composition of the vapour phase in equilibrium with the solution is determined by the partial pressure of the components. If y₁ and y₂ are the mole fractions of the components 1 and 2 in the vapour phase, then, using Dalton’s law of partial pressures
P₁ = y₁ Ptotal
P₂ = y₂ Ptotal
In general P_i = y_i P_total

It can be concluded: At the equilibrium vapour phase will be always richer in the component which is more volatile.
Raoult’s law is a special case of Henry’s law

According to Raoult’s law p_i = x_i pi°
According to Henry’s Law p = K_H × x

Here only the proportionality constant K_H differs from p_i°. Thus Raoult’s law becomes a special case of Henry’s law in which K_H becomes equal to P_i°.

Ideal and Non-Ideal Solution: A solution that obeys Raoult’s Law at all concentrations and at all temperatures is called an Ideal Solution. Herein the magnitude of solute-solvent interactions is the same as the magnitude and solvent of the solute-solute interaction- solvent interactions magnitude in the two components.

For ideal solutions enthalpy on mixing remains the same,
i.e., ΔH_Mix = 0

For such solutions, there is no change in volume on mixing the two components
i.e., ΔV_Mix = 0

Non-Ideal Solutions are those which do not obey Raoult’s Law: For such solutions, the magnitude of solute-solvent interactions is either greater than or less than the magnitude of solute-solute or solvent-solvent interactions in the pure components.

Additionally, during solution formation either heat is evolved or absorbed.
i.e.; ΔH_Mix = +Ve
or
ΔH_Mix = -Ve [or ΔH_Mix ≠ 0]
Moreover, there is a volume change in mixing
ΔH_Mix ≠ 0

Examples of ideal solutions:

1. n-hexane + n-heptane
2. Benzene + Toluene
3. Chlorbenzene + bromobenzene
4. CCl₄ + SiCl₄
5. CH₃OH + C₂H₅OH.

Examples of Non-ideal solutions:

With Positive deviations	With Negative deviations
1. acetone and ethanol	1. Water + HCl
2. Acetone + Benzene	2. Water + HNO3
3. Acetone + CCl4	3. CHCl3 and acetone
4. Benzene + CCl4	4. Acetic acid + pyridine
5. Water + ethyl alcohol	5. Acetone + aniline
6. Cyclohexane + ethanol	6. Benzene + CHCl3

Colligative Properties: These properties depend only on the number of solute particles, but not on the nature of the solute.
They are

1. Relative lowering of vapour pressure
2. Elevation in boiling points
3. Depression in freezing point
4. Osmotic pressure.

1. Relative lowering of vapour pressure: The relative lowering of the vapour pressure of a dilute solution (which, for all intents and purposes, is an ideal solution) is equal to the mole fraction of the solute at a given temperature.
If x_A = Mole fraction of the solvent
x_B = Mole fraction of the solute
P_A = Vapour pressure of the pure solvent
P_B = Vapour pressure of the solution

\(\frac{P_{A}^{\circ}-P_{A}}{P_{A}^{\circ}}\) is called the relative lowering of vapour pressure
W_A = Mass of comp. A
M_A = M. mass of comp. A
W_B = Mass of comp. B
M_B = M. mass of comp. B

From relative lowering in vapour pressure, the molar mass of solute can be calculated as:
M_B = \(\frac{W_{\mathrm{B}} \times \mathrm{M}_{\mathrm{A}}}{\mathrm{W}_{\mathrm{A}} \times \frac{\mathrm{P}_{\mathrm{A}}^{\circ}-\mathrm{P}_{\mathrm{A}}}{\mathrm{P}_{\mathrm{A}}^{\circ}}}\)

2. Elevation in boiling: The boiling point of a liquid may be defined as the temperature at which its vapour pressure becomes equal to atmospheric pressure. When a non-volatile solute is added to a solvent, the solution boils at a higher temperature than the pure solvent.

If T_b° is the boiling point of the pure solvent and T_b is the boiling point of the solution, then elevation in boiling point,
ΔT_b = T_b – T_b° = K_b × m
where m is the molality of the solution and K_b is molal elevation constant or ebullioscopic constant of the solvent.

Molal elevation in boiling point is defined as the elevation in boiling point produced when 1 mole of a solute is dissolved in 1000 g of the solvent.

The molar mass of the solute can be calculated from the elevation in boiling point as
M_B = \(\frac{\mathrm{K}_{\mathrm{b}} \times \mathrm{W}_{\mathrm{B}} \times 1000}{\Delta \mathrm{T}_{\mathrm{b}} \times \mathrm{W}_{\mathrm{A}}}\)

The molar elevation constant is expressed as
K_b = \(\frac{\mathrm{RT}_{\mathrm{b}}^{\circ} \mathrm{M}}{\Delta \mathrm{H}_{\mathrm{vap}} \times 1000}\)
or
= \(\frac{\mathrm{RT}_{\mathrm{b}}^{\circ}}{\Delta \mathrm{h}_{\mathrm{vap}} \times 1000}\)

where T_b° is the boiling point of the solvent, M is the molar mass of the solvent, R is gas constant and ΔH_vap is the enthalpy of vaporisation of the pure solvent. Δh_vap is the enthalpy of vaporisation per gram as

Δh_vap = \(\frac{\Delta \mathrm{H}_{\mathrm{vap}}}{\mathrm{M}}\)

The vapour pressure curve of the solution lies below the curve for pure solvent. The diagram shows that ΔT_b denotes the elevation of the boiling point of a solvent in the solution.

3. Depression in freezing point: The freezing point of a liquid may be defined as the temperature at which solid and the liquid states of the same substance have the same vapour pressure. When a non-volatile solute is added to a solvent, the freezing point of the solution is always less than the freezing point of the pure solvent.

If T_f° is the freezing point of the pure solvent and T_f is the freezing point of the solution, then depression in freezing point,
ΔT_f = T_f° – T_f = K_f. m
where K_f = molal depression constant or cryoscopic constant
m = molality of the solution

Molal depression constant (K_f) is the depression in freezing point produced when 1 mole of solute is dissolved in 1000 g of the solvent.

Molar mass of solute from depression in freezing point may be calculated as
M_B = \(\frac{\mathrm{K}_{\mathrm{f}} \times \mathrm{W}_{\mathrm{B}} \times 1000}{\Delta \mathrm{T}_{\mathrm{f}} \times \mathrm{W}_{\mathrm{A}}}\)

The molal depression constant is
K_f = \(\frac{\mathrm{RT}_{\mathrm{f}}^{\circ} \mathrm{M}}{\Delta \mathrm{H}_{\mathrm{t}} \times 1000}\)
or
= \(\frac{R T_{f}^{\circ} f}{\Delta h_{f} \times 1000}\)
where R is gas constant, Tf° is the freezing point of the solvent, M is the molar mass and ΔHf is the enthalpy of fusion of the solvent, Δhf is the enthalpy of fusion per gram and is

Δh_f = \(\frac{\Delta \mathrm{H}_{\mathrm{f}}}{\mathrm{M}}\)

Diagram showing ΔT_f depression of the freezing point of a solvent in a solution.

4. Osmotic pressure: The phenomenon of the flow of solvent through a semi-permeable membrane from pure solvent to the solution ‘ is called osmosis. The osmotic pressure may be defined as the excess pressure which must be applied to a solution to prevent the passage of solvent into it through a semi-permeable membrane. The osmotic pressure for a solution (TC) is
π = c RT
where c is the molar concentration, R is the gas constant and T is the temperature.

Molar mass of a solute from the osmotic pressure may be calculated as:
MB = \(\frac{W_{B} R T}{V \times \pi}\)

→ Isotonic solutions: The solutions having the same osmotic pressure are called isotonic.

→ Abnormal Molar Masses: Colligative properties depend upon the number of solute particles. Therefore, in case of those solutions which undergo dissociation or association in solution, the colligative properties will be different
Van t Hoff factor = \(\frac{\text { Observed colligative property }}{\text { Normal value of colligative property }}\)
= \(\frac{\pi_{\text {obs }}}{\pi_{\text {nor }}}=\frac{\Delta \mathrm{T}_{\mathrm{b}(\mathrm{obs})}}{\Delta \mathrm{T}_{\mathrm{b}(\mathrm{nor})}}=\frac{\Delta \mathrm{T}_{\mathrm{f}(\mathrm{obs})}}{\Delta \mathrm{T}_{\mathrm{f}(\text { nor })}}\)

→ The principle of measuring osmotic pressure: The pressure in excess of atmospheric pressure that must be applied to the solution to prevent it from rising in the tube is the osmotic pressure. This will also be equal to”hydrostatic pressure of the liquid column of height, h.

Since colligative property is inversely proportional to the molar mass of the solute, therefore,
i = \(\frac{\text { Normal molar mass }}{\text { Observed molar mass }}\)

∴ Observed Osmotic pressure, π = i × \(\frac{n}{V}\) RT
Observed elevation in B.Pt. ΔT_b = i × K_b × m
Observed depression in F.Pt. ΔT_f = i × K_f × m where m = molality of the solution

Reverse Osmosis and water purification: The direction of Osmosis can be reversed if a pressure larger than the osmotic pressure is applied to the solution, side. That is, now the pure solvent flows out of the solution tiìrough the semi-permeable membrane (SPM).

This phenomenon is called Reverse Osmosis and is of great utility. Reverse osmosis is used for the Desalination of seawater. The pressure required for reverse osmosis is quite high and a workable porous membrane is a film of cellulose acetate placed over suitable support. Cellulose acetate is permeable to water but impermeable to impurities and ions present in seawater.

Reverse osmosis occurs when a pressure larger than the osmotic pressure is applied to the solution.